THE 

ELECTROLYTIC   DISSOCIATION  THEORY 
WITH  SOME  OF  ITS  APPLICATIONS 


AN  ELEMENTARY  TREATISE 

FOR  THE  USE  OF 
STUDENTS  OF  CHEMISTRY 


BY 

HENRY  P.  TALBOT,  PH.D. 
PROFESSOR  OF   INORGANIC  AND  ANALYTICAL  CHEMISTRY 

AND 

ARTHUR  A.  BLANCHARD,  PH.D. 

INSTRUCTOR   IN   INORGANIC  CHEMISTRY 

AT  THE   MASSACHUSETTS   INSTITUTE  OF  TECHNOLOGY 


OF 


NEW    YORK: 
THE    MACMILLAN    CO. 

LONDON  :    MACMILLAN  &  Co.,  Ltd. 


COPYRIGHT  1905 
HENRY  P.  TALBOT 

AND 

ARTHUR  A.    BLANCHARD 


THOMAS    TODD,  PRINTER 
J4   BEACON    STREET,  BOSTON,  MASS. 


PREFACE 

IN  this  little  treatise  the  authors  have  sought  to  bring  together,  in  small 
compass,  material  relating  to  the  Electrolytic  Dissociation  Theory  which  is 
now  somewhat  widely  distributed  throughout  many  of  our  excellent  text- 
books. The  method  of  presentation  is  that  which  they  have  found  service- 
able in  enabling  their  students  to  comprehend  the  main  facts  which  are  today 
generally  accepted  as  supporting  the  Theory,  and  to  understand  its  application 
to  important  types  of  chemical  change. 

While  the  book  has  been  written  primarily  for  the  use  of  students,  the 
authors  have  also  kept  in  mind  its  probable  usefulness  to  the  teacher  in 
preparatory  school  or  college,  who  may  desire  to  gain  in  a  short  time  an 
acquaintance  with  the  fundamental  facts  and  principles  in  this  interesting 
field.  In  this  connection,  however,  they  desire  to  express  their  firm  convic- 
tion that  this  Theory  should  be  touched  upon  in  only  the  most  elementary 
way  in  the  secondary  schools;  but  this  does  not,  of  course,  make  it  less 
necessary  that  the  well-informed  teacher  should  be  prepared  to  meet  the 
inquiries  of  the  occasional  unusually  mature  and  thoughtful  pupil. 

More  has  been  included  in  this  manual  than  the  authors  have  found  it 
advisable  for  the  college  student  who  is  just  beginning  the  study  of  chemical 
science  to  attempt  to  master,  and  quite  as  much  as  many  students  who  have 
already  had  a  year  of  chemical  experience  in  a  preparatory  school  will  be 
able  to  thoroughly  understand.  Our  experience  has  shown,  however,  that 
it  is  easier  to  maintain  the  interest  of  the  thoughtful  pupil  if  answers  to  some 
of  the  questions  suggested  to  him  by  the  discussions  in  the  main  body  of  the 
text  are  placed  at  his  hand,  when  it  is  possible  to  do  this  without  going  too 
far  afield.  Most  of  the  material  of  this  nature  has  been  printed  in  smaller 
type,  and  may  be  omitted  without  loss  of  continuity ;  and  it  is  believed  that 
in  the  course  of  a  year  of  study  the  college  student  (even  the  beginner  in 
chemistry,  if  he  has  some  knowledge  of  physics)  can  be  brought  to  under- 
stand the  essential  principles  included  in  the  main  text.  To  insure  this  it 
is  necessary  that  the  instructor  should  lose  no  opportunity  throughout  his 
course  to  emphasize  the  application  of  a  principle,  after  it  has  once  been 
introduced. 


iv  Preface 

The  attempt  to  present  the  subject-matter  in  a  simple  form,  and  at  the 
same  time  to  avoid  inaccurate  statements,  has  sometimes  led  to  a  conflict  of 
ideals,  as,  for  example,  in  the  application  of  the  Law  of  Mass  Action  to 
strong  electrolytes.  Nothing  is  said  in  this  connection  of  the  unexplained 
fact  that  such  electrolytes  apparently  do  not  rigidly  obey  this  law,  since  to 
discuss  this  topic  would  seriously  complicate  an  important  statement  without 
corresponding  advantage.  We  believe,  however,  that  in  such  cases  no  vio- 
lence has  been  done  to  the  principles  of  physical  chemistry,  and  that  no 
impressions  have  been  given  which  it  will  be  difficult  for  the  student  to 
unlearn  if  he  pursues  the  subject  to  its  more  advanced  stages. 

The  application  of  the  ionic  theory  to  indicators  has  been  omitted  alto- 
gether, as  the  present  state  of  our  knowledge  seems  to  indicate  that  compli- 
cated rearrangements  of  the  atoms  within  the  molecules  of  organic  substances 
are  involved,  with  which  it  would  be  beyond  the  scope  of  this  treatise  to  deal. 

The  authors  desire  to  acknowledge  their  indebtedness  to  many  of  the 
standard  text-books,  and  especially  to  Smith's  "  Laboratory  Outline  of  Gen- 
eral Chemistry  "  for  suggestions  as  to  laboratory  experiments.  They  would 
also  acknowledge  the  valuable  assistance  rendered  by  Dr.  Miles  S.  Sherrill, 
and  the  friendly  and  helpful  criticisms  of  other  members  of  the  instructing 
staff  of  the  Massachusetts  Institute  of  Technology. 

H.  P.  TALBOT. 

A.  A.  BLANCHARD. 

September,  1905. 


CONTENTS 

CHAPTER  I.     EVIDENCES  OF  ELECTROLYTIC   DISSOCIATION   AFFORDED   BY   A 

STUDY  OF  THE  PROPERTIES  OF  SOLUTIONS        ....  i 

The  Lowering  of  the  Freezing  Point           ........  2 

The  Raising  of  the  Boiling  Point 5 

Osmotic  Pressure         ............  6 

Chemical  Activity 10 

Electrolytic  Conduction 1 1 

Polarity  in  Chemical  Compounds II 

Ions 12 

Movement  of  the  Ions 14 

Electrical  Charges  upon  the  Ions     .         .         .         .         .         .         .         .  16 

Reactions  at  the  Electrodes 17 

Summary 19 

CHAPTER  II.     THE  LAW  OF  MASS  ACTION  AND  THE  CHEMICAL  BEHAVIOR  OF 

ELECTROLYTES    21 

Reversible  Reactions  and  the  Effect  of  Mass     .         .         .         .         .         .         .  21 

Degree  of  lonization  ............  25 

The  Formation  of  Insoluble  Compounds  through  the  Interaction  of  Certain 

Ions 26 

Solubility  Product 28 

Characteristic  Reactions  of  the  Various  Ions 31 

Simple  Ions 31 

Complex  Ions 32 

Acids 33 

Bases        .                  34 

Salts 34 

Neutralization 35 

The  Action  of  a  Strong  Acid  upon  the  Salt  of  a  Weak  Acid     ....  37 

The  Effect  of  the  Formation  of  Volatile  Products           ....  38 

The  Solubility  of  Carbonates  in  Acids    .......  39 

The  Precipitation  of  the  Metallic  Sulphides .40 

The  Action  of  a  Strong  Base  upon  the  Salt  of  a  Weak  Base  .         .         .         .  41 
Effect  upon  the  Properties  of  Weak  Acids  or  Bases  of  Neutral  Salts  with 

a  Common  Ion 41 

Hydrolysis  ..............  42 

CHAPTER  III.     ELECTROLYTIC  SOLUTION  PRESSURE 45 

The  Electrolytic  Solution  Pressure  of  the  Metals 45 

The  Electrolytic  lonization  Pressure  of  the  Negative  Elements,  or  the  Ten- 
dency of  these  Elements  to  Pass  into  the  Ionic  Condition     ....  49 

CHAPTER  IV.     OXIDATION  AND  REDUCTION         ...        .        .        .        .  51 

CHAPTER  V.     THE  MORE  COMMON  IONS  AND  THEIR  CHARACTERISTICS  .        .  57 

The  Ions  of  the  Metals       .         .                 -. 58 

The  Ions  of  the  Non-metals        . 67 

CHAPTER  VI.     EXPERIMENTS          .        .        .  '     .        .        .        .        .        .        .  71 

APPENDIX          .'      . ;        .      -.        .  81 

Degree  of  Dissociation  of  Some  of  the  Most  Important  Electrolytes         .         .  81 

v 


CAUfgS^ 


CHAPTER    I 

EVIDENCES    OF    ELECTROLYTIC    DISSOCIATION    AFFORDED    BY    A 
STUDY    OF    THE    PROPERTIES    OF    SOLUTIONS 

i.  If  the  various  forms  of  matter  are  studied  with  reference  to 
their  ability  to  transmit  an  electric  current,  it  is  found  that  some,  such 
as  glass,  hard  rubber,  or  alcohol,  do  not  allow  an  appreciable  amount 
of  electricity  to  pass  through  them,  while  others  permit  the  passage  of 
electricity  with  comparative  readiness.  The  former  are  called  non- 
conductors ;  the  latter,  conductors.  The  conductors,  in  turn,  may  be 
subdivided  into  two  classes  with  reference  to  the  manner  in  which 
they  transmit  the  current,  namely,  metallic  conductors  and  electrolytic 
conductors.  Platinum  or  copper  wires,  or  the  carbon  filaments  of  incan- 
descent lamps,  are  of  the  first  variety,  and  conduct  the  current  without 
undergoing  any  permanent  alterations,  while  in  conductors  of  the  sec- 
ond variety  (comprising  salts  when  in  the  molten  condition,  or  solutions, 
particularly  aqueous  solutions,  of  acids,  bases,  or  salts)  the  passage  of 
the  current  is  accompanied  by  a  separation  of  the  components  of  the 
conductor.  This  electrolytic  conduction  can  only  take  place  in  fluid 
bodies  the  components  of  which  gradually  collect  at  the  poles,  that  is, 
at  the  points  where  the  current  enters  and  leaves  the  fluid.  Bodies 
which  conduct  in  this  manner  are  known  as  electrolytes,  and  the  proc- 
ess of  conduction  accompanied  by  the  separation  of  the  constituents 
of  the  electrolyte  is  called  electrolysis. 

It  was  to  explain  electrolysis,  together  with  certain  other  peculiar 
properties  possessed  only  by  those  solutions  which  conduct  electricity, 
that  in  1887  Arrhenius,  a  Swedish  physicist,  was  led  to  propose  the 
Theory  of  Electrolytic  Dissociation. 

This  theory  assumes  that  a  certain  proportion  of  the  dissolved  mole- 
cules of  electrolytes  are  dissociated  into  simpler  component  parts,  called 
ions,  and  that  each  of  these  ions  has  an  effect  upon  many  of  the  proper- 
ties of  solutions  equal  to  that  of  a  whole  molecule.  It  is  the  purpose 


2  Electrolytic  Dissociation   Theory 

of  the  following  pages  to  present  some  of  the  important  considerations 
which  have  led  to  the  general  acceptance  of  this  theory,  and  to  illustrate 
the  applications  of  the  theory  to  certain  typical  chemical  changes. 

With  a  very  few  exceptions  all  solid  non-metallic  substances,  as  well 
as  all  pure  liquids  and  gases,  are  non-conductors  of  electricity  at  ordinary 
temperatures.  Thus  a  crystal  of  sodium  chloride  is  a  non-conductor, 
and  pure  water  is  a  non-conductor.  If,  however,  the  crystallized  sodium 
chloride  is  heated  until  it  fuses,  or,  still  more  easily,  if  it  is  dissolved  in 
water,  the  result  is  an  excellent  conductor  of  electricity.  The  discus- 
sion which  follows  will  be  confined  mainly  to  the  latter  case,  namely, 
that  in  which  a  non-conductor,  upon  being  dissolved  in  a  non-conducting 
liquid  at  ordinary  temperatures,  becomes  a  good  conductor  of  electricity. 

It  must  first  be  noted  that  not  all  solutions  exhibit  this  phenomenon 
of  electrolytic  conduction,  and  that  those  solutions  which  do  possess 
this  property  also  possess  certain  other  remarkable  properties  which  are 
absent  in  solutions  which  do  not  allow  a  current  to  pass  through  them. 
These  properties  may  be  summarized  as  follows  :  — 

1.  Such  solutions  freeze  at  an  abnormally  low  temperature. 
II.     They  boil  at  an  abnormally  high  temperature. 

III.  The  osmotic  pressure  of  the  dissolved  substance  is  abnormally  great. 

IV.  The  dissolved  substance  exhibits  great  chemical  activity. 
V.     They  permit  the  passage  of  an  electric  current. 

These  five  striking  properties  will  first  be  considered,  and  it  will 
not  only  be  seen  that  if  each  is  carefully  examined  it  leads  to  the  con- 
clusion expressed  by  the  Dissociation  Theory,  as  outlined  above,  but 
it  will  be  further  shown  that  if  the  proportion  of  the  molecules  of  an 
electrolyte  which  are  dissociated  is  estimated  from  data  based  upon 
an  independent  study  of  each  of  the  properties,  the  values  obtained  are 
substantially  concordant. 

THE  LOWERING  OF  THE  FREEZING  POINT 

2.  When  any  substance  is  dissolved  in  a  pure  liquid  the  resulting 
solution  freezes  at  a  lower  temperature  than  the  solvent  alone,  and  for 
any  given  substance  the  amount  by  which  the  freezing  point  is  lowered 
below  that  of  the  pure  solvent  is  in  proportion  to  the  quantity  of  the 


Lowering  of  the  Freezing  Point  3 

substance  dissolved ;  for  example,  the  lowering  of  the  freezing  point 
by  the  solution  of  20  grams  of  sugar  in  a  specific  volume  of  water 
(say  1,000  c.c.)  is  twice  as  great  as  that  produced  by  the  solution  of 
10  grams  in  the  same  volume.  Moreover,  if  60  grams  of  urea,  46 
grams  of  ethyl  alcohol,  and  342  grams  of  sugar  are  each  dissolved  in 
1,000  grams  of  water,  the  resulting  solutions  freeze  at  —  i.86°C. 
The  quantities  named  are  the  molecular  weights  in  grams  respectively 
of  the  three  substances,  and  these  amounts  will  hereafter  be  designated 
as  mols. 

The  mol  (60  grams)  of  urea  is,  of  course,  made  up  of  an  exceedingly 
large  number  of  the  very  minute  individual  molecules,  and  the  same  is  true 
of  the  mol  (342  grams)  of  sugar;  yet,  since  the  quantities  expressed  by  the 
mols  stand  in  the  same  relation  to  each  other  as  the  molecular  weights  of 
the  two  substances,  the  number  of  actual  molecules  must  be  the  same  in 
both  cases.  Solutions  which  contain,  in  a  given  volume,  any  other  quan- 
tities of  these  substances  which  stand  in  the  same  proportion  to  each  other 
as  their  molecular  weights  must  also  contain  the  same  actual  number  of 
molecules ;  that  is,  be  equi-molal. 

The  above  statement  holds  true  for  all  bodies  which  do  not  yield 
solutions  which  conduct  an  electric  current.  Such  bodies  are  termed 
non-electrolytes,  in  distinction  from  those  bodies  which  in  solution  do 
conduct  a  current,  and  are  termed  electrolytes.  Thus,  the  freezing 
point  of  a  solution  of  any  non-electrolyte  containing  i  mol  in  1,000 
grams  of  water  is  1 .86°  C.  lower  than  the  freezing  point  of  the  water. 
This  quantity  is  known  as  the  molecular  lowering  of  the  freezing  point, 
and  it  is  easy  to  see  that  if  this  molecular  lowering  for  water  has  once 
been  established  from  experiments  with  a  considerable  number  of  sub- 
stances of  known  molecular  weight,  then  the  number  of  mols  of  any 
non-electrolyte  dissolved  in  water  may  be  found  by  determining  the 
freezing  point  of  its  solution. 

For  example,  a  solution  of  50  grams  of  wood  alcohol  in  1,000 
grams  of  water  is  found  to  freeze  at  —2.90°.  Since  every  mol  of 
a  non-electrolyte  should  cause  a  lowering  of  the  freezing  point  of  1 .86°, 

2  QO 

the  number  of  mols  in  the  solution  must  be  -  ~,  or  1.56.     This  is  in 

I.oO 

accordance  with  the  facts,  because  the  amount  dissolved  (50  grams) 
must  represent  — ,  or  1.56  mols,  since  the  mol  of  wood  alcohol 
(CH3OH)  is  32  grams. 


4  Electrolytic  Dissociation   Theory 

If  the  lowering  of  the  freezing  point  produced  by  a  given  number 
of  mols  of  an  electrolyte  is  compared  with  that  produced  by  the  same 
number  of  mols  of  a  non-electrolyte  (such  as  the  wood  alcohol  of  the 
preceding  paragraph)  it  is  found  to  be  abnormally  large.  If,  however, 
the  principle  stated  above,  that  the  number  of  mols  of  the  dissolved 
substance  may  be  measured  by  the  extent  of  the  lowering  of  the 
freezing  point,  holds  true,  then  it  must  follow  that  a  given  number 
of  molecules  of  an  electrolyte  (using  the  term  molecule  in  its  ordi- 
nary chemical  sense),  when  in  solution,  are  in  reality  broken  up  into 
a  greater  number  of  smaller  molecules  of  some  sort,  each  one  of  which 
has  the  same  effect  in  depressing  the  freezing  point  of  the  solution  as 
one  of  the  original  undecomposed  molecules,  or  as  a  molecule  of  a 
non-electrolyte. 

An  example  will  make  this  clear.  A  solution  containing  5.85  grams 
of  sodium  chloride  to  the  liter  is  found  by  experiment  to  freeze  at 
—  0.350°  C.  If  sodium  chloride  were  a  non-electrolyte,  that  is,  if  it 
formed  only  NaCl-molecules  in  a  solution,  the  freezing  point  would 
be  —  o.i  86°,  because,  since  this  solution  contains  5.85  grams  of  dis- 
solved substance  in  1,000  grams  of  solvent,  and  the  molecular  weight 
of  NaCl  is  58.5,  the  5.8$  grams  should  constitute  o.ioo  mol,  and 

c  8  c 
should  depress  the  freezing  point   to  the  extent  of  ^5-^  X  1.86,   or 


0.186°.      That   the  actual    lowering  is  0.350°,  or  =  1.88  times 

.100 

greater  than  that  which  should  result  if  the  sodium  chloride  behaved 
in  the  same  manner  as  the  non-electrolytes  illustrated  above,  shows 
that  the  solution  contains  a  larger  number  of  molecules  than  the  num- 
ber of  chemical  molecules  (NaCl).  It  is  easily  seen  that  this  would 
be  explained  if,  for  example,  out  of  1,000  of  the  original  NaCl-mole- 
cules  880  should,  when  dissolved,  dissociate  into  the  smaller  individuals 
Na  and  Cl,  while  120  remained  in  the  form  of  the  larger  NaCl  aggre- 
gates. The  total  number  of  molecules  would  then  be  1,880,  which  is 
1.88  times  the  number  of  the  chemical  molecules  (NaCl.)  Likewise, 
of  1,000  molecules  of  such  a  body  as  calcium  chloride,  CaCl2,  750 
molecules  might  dissociate  into  750  of  the  smaller  individuals  Ca  and 
1,500  of  Cl,  while  there  remained  250  of  the  CaCl2  aggregates,  thus 


Raising  of  the  Boiling  Point  5 

giving  2,500  molecules  in  solution.  This  supposition  is  in  accord 
with  the  actual  freezing  point  of  a  calcium  chloride  solution,  which 
indicates  2.5  times  as  many  molecules  as  the  number  of  chemical 
molecules  (CaCl2). 

It  may  at  once  be  asked  whether  such  a  supposition  as  is  here 
suggested  is  justifiable,  since  a  solution  of  sodium  chloride  exhibits 
none  of  the  well-known  properties  of  sodium  or  chlorine.  The  full 
•explanation  of  the  difference  between  these  smaller  molecules  (part- 
molecules)  and  ordinary  atoms  of  elements  in  the  uncombined  state 
can  best  be  given  in  a  later  paragraph  (see  page  12),  but  it  may  be 
briefly  stated  here  that  it  is  due  to  the  electrical  charges  which  reside 
upon  these  part-molecules,  essentially  altering  their  behavior.  In  such 
bodies  as  ammonium  nitrate,  the  breaking  down  of  the  chemical  mole- 
cules is  not  into  single  atoms,  but  into  groups  of  atoms,  as  NH4  and 
NO3,  which  also  bear  electrical  charges. 

THE  RAISING  OF  THE  BOILING  POINT 

3.  If  any  non-volatile  substance  is  dissolved  in  any  liquid,  the 
temperature  at  which  that  liquid  boils  is  raised.  The  same  principles 
are  found  to  hold  with  respect  to  the  raising  of  the  boiling  point  which 
have  just  been  outlined  for  the  lowering  of  the  freezing  point.  This 
property  of  solutions  may,  therefore,  likewise  be  made  use  of  to  deter- 
mine the  number  of  molecules  of  the  dissolved  substance  which  a 
solution  contains. 

The  molecular  raising  of  the  boiling  point  of  water  is  0.52°  ;  that  is, 
if  i  liter  of  an  aqueous  solution  contains  i  mol  of  any  non-electro- 
lyte which  does  not  itself  boil  at  a  low  temperature,  the  solution  boils 
0.52°  C.  higher  than  pure  water.  All  solutions  of  electrolytes  boil  at 
abnormally  high  temperatures  just  as  they  freeze  at  abnormally  low 
temperatures,  and  the  extent  to  which  each  electrolyte  is  dissociated 
into  smaller  molecules  is  shown  to  be  the  same  by  means  of  this  method 
as  by  the  method  previously  discussed.  For  example,  a  3.3  per  cent. 
solution  of  sodium  chloride  boils  at  100.50°  C.  If  it  were  an  undis- 
sociated  body  the  boiling  point  would  be  raised  only  to  the  extent  of 

°-52»  or  °-3°°-     Tnat  tne  actual  raising  is    —  —  »    or  1.7  times 


-»         --  —  — 

50.5  0.30 


6  Electrolytic  Dissociation   Theory 

this  amount,  leads  to  the  same  conclusion  as  before,  namely,  that  in. 
solution  about  0.7  of  the  NaCl  aggregates  are  dissociated  into  smaller 
molecules.  (The  difference  between  this  value  and  that  deduced  from 
the  freezing  point  lowering  [0.88]  is  not  greater  than  would  result 
from  differing  conditions  and  experimental  difficulties.) 

OSMOTIC  PRESSURE 

4.  An  understanding  of  the  nature  of  the  phenomenon  known  as 
osmotic  pressure  can  best  be  gained  through  a  consideration  of  gas 
pressure,  which  it  closely  resembles. 

The  molecules  of  a  gas,  as  indeed  of  any  substance  not  at  the 
absolute  zero  of  temperature,  are  in  a  state  of  rapid  vibration,  and,  in 
their  movement  to  and  fro,  the  distance  which  a  molecule  may  move 
in  a  straight  line  is  limited  to  the  space  through  which  it  may  travel 
without  colliding  with  another  molecule,  or  with  the  walls  of  the  con- 
taining vessel.  When  any  two  molecules  of  a  gas  collide  they  at  once 
rebound  with  such  force  that  each  immediately  escapes  beyond  the 
other's  sphere  of  attraction,  while  the  impacts  of  the  molecules  upon 
the  walls  of  the  containing  vessel,  against  which  they  strike  and  then 
rebound,  give  rise  to  the  phenomenon  of  gas  pressure. 

If  the  volume  of  a  given  quantity  of  gas  is  increased  without 
changing  the  temperature,  its  pressure  is  diminished.  This  must  fol- 
low if  the  statement  of  the  nature  of  gas  pressure  is  correct,  because 
by  increasing  the  volume  the  number  of  impacts  which  can  be  made 
by  the  molecules  of  the  gas  on  a  given  area  of  containing  wall  is 
decreased.  This  is  in  accordance  with  Boyle's  Law,  which  states  that 
for  a  given  quantity  of  gas  kept  at  a  constant  temperature  the  pressure 
is  inversely  proportional  to  the  volume,  or  that 

pressure  X  volume  =  constant. 

If  the  temperature  of  a  gas  is  increased  without  allowing  its  volume 
to  change,  its  pressure  is  increased.  This  follows  from  the  fact  that 
temperature  is  the  measure  of  the  intensity  of  the  vibratory  motion.. of 
the  molecules  of  a  substance.  If  the  molecules  of  a  gas  vibrate  faster, 
they  must  necessarily  strike  oftener  and  harder  upon  the  walls  of  the 
containing  vessel,  or,  in  other  words,  they  must  exert  a  greater  pres- 


Osmotic  Pressure  7 

sure.  This  accords  with  the  Law  of  Charles,  which  states  that  the 
pressure  of  a  gas,  kept  at  a  constant  volume,  is  proportional  to  the 
absolute  temperature. 

If  two  vessels  of  the  same  size  contain  the  same  number  of  mole- 
cules of  different  gases  at  the  same  temperature,  each  gas  exerts  the 
same  pressure.  This  is  the  converse  of  Avogadro's  Theory,  which 
states  that  two  equal  volumes  of  gas  at  the  same  temperature  and 
pressure  contain  the  same  number  of  molecules.  From  this  it  appears 
that  the  pressure  exerted  by  a  body  of  gas  at  a  given  temperature  and 
volume  depends  solely  upon  the  number,  and  in  no  way  upon  the  kind, 
of  its  molecules.  In  other  words,  the  rapidity  of  vibration  of  the  mole- 
cules of  hydrogen  at  any  given  temperature  so  far  exceeds  that  of  the 
molecules  of  oxygen  that  a  given  number  of  the  molecules  of  the  former 
element  exerts  the  same  pressure  as  the  same  number  of  molecules  of 
the  latter  element.  The  pressures  of  two  equal  volumes  of  gases  at  the 
same  temperature  stand,  therefore,  in  the  same  relations  as  the  numbers 
of  molecules  of  the  gases  in  these  volumes. 

A  molecular  weight  in  grams  (mol)  of  any  gas,  as  17  grams  of 
ammonia  gas  (NH3)  or  2  grams  of  hydrogen,  at  the  standard  pressure, 
760  mm.  of  mercury,  and  at  the  standard  temperature,  o°  C.,  occupies 
a  volume  equal  to  22.4  liters;  if  it  is  compressed  to  i  liter,  it  will 
exert  a  pressure  equal  to  22.4  atmospheres,  i.  c.,  22.4  X  760  mm. 

5.  When  a  substance  is  dissolved  its  molecules  vibrate  to  and  fro 
in  the  solvent  in  the  same  manner  as  do  the  molecules  of  a  gas  in 
the  space  which  the  gas  occupies.  These  molecules  of  dissolved  sub- 
stance will,  therefore,  exert  a  pressure  against  an  opposing  surface  in 
exactly  the  same  manner  as  the  molecules  of  a  gas.  This  is  known 
as  osmotic  pressure.  If  by  any  means  it  were  possible  to  measure  this 
osmotic  pressure  of  the  dissolved  molecules  separately  from  the  pres- 
sure of  the  solvent,  it  is  evident  from  what  has  been  said  concerning 
gas  pressure  that  such  measurements  might  furnish  another  means  of 
determining  the  number  of  molecules  of  the  substance  in  solution.  To 
illustrate  how  such  measurements  might  be  possible,  let  us  imagine  a 
fish  seine  with  which  fishermen  have  succeeded  in  encircling  a  school 
of  mackerel.  Water  passes  freely  through  the  net  and  exerts  no  pres- 
sure, but  the  fish  are  too  large  to  pass  through  the  meshes.  In  their 


Electrolytic  Dissociation   Theory 

efforts  to  escape  they  strike  against  the  net  and  thus  exert  a  pressure 
upon  it.  Similarly,  if  a  mesh  can  be  prepared  so  fine  that  molecules  of 
a  substance  dissolved  in  water  cannot  pass  through  it,  but  still  so  large 
that  water  molecules  may  readily  pass,  it  would  furnish  a  means  of 
measuring  osmotic  pressure.  Such  meshes  do  exist,  and  are  called 
osmotic  membranes.  The  walls  of  plant  cells,  many  animal  mem- 
branes, and  many  artificially  prepared  films  are  of  this  semi-permeable 
character. 

6.  The   formation   of   such   membranes,    as   well   as   the   existence   of 
osmotic  pressure,  may  be  qualitatively  shown  by  what  may  be  called  the 
mineral floiver  garden,  prepared  as  follows:  small  lumps  or  crystals  of  cer- 
tain very  soluble  salts,  e.g.,  ferric  chloride,  copper  chloride,  nickel  nitrate, 
cobalt  chloride,  or  manganese  sulphate,  are  dropped  into  a  solution  of  sodium 
silicate  (water  glass,  sp.  gr.  i.i).     Their  behavior  resembles  that  of  growing 
seeds,  as  they  appear  to  immediately  sprout  and  send  up  shoots  toward  the 
surface  of  the  liquid,  which  grow  with  a  visible  rapidity.     In  fact,  the  salts 
have  at  once  commenced  to  dissolve,  forming  thin  layers  of  very  concentrated 
solution  about  each  lump.     At  the  surface,  separating  each  of  these  layers 
of  solution  from  the  water  glass,  there  forms  a  film  of  the  insoluble  silicate  of 
the  metal.     This  film  is  an  osmotic  membrane  which  allows  water  to  pass 
either  in  or  out,  but  the  molecules  of  salt,  not  being  able  to  pass  through, 
exert  against  it  their  osmotic  pressure,  and  break  it  at  its  weakest  part,  which 
is  always  the  top.     This  exposes  a  new  surface  of  the  salt  solution  to  the 
sodium  silicate,  and  a  new  film  forms,  which,  in  turn,  is  broken,  thus  permit- 
ting the  growth  of  the  little  tube  of  the  silicate  of  the  metal.     Clusters  of 
these  tubes  of  various  colors  give  an  appearance  of  plant  growth  within  the 
liquid. 

7.  The  quantitative  measurement  of  osmotic  pressure,  although  a 
matter  of  considerable  difficulty,  may  still  be  made  with  a  very  fair 
degree  of  accuracy.     A  solution  of  342  grams  of  sugar  in  22.4  liters  of 
water  at  o°  C.  gives  an  osmotic  pressure  of  I  atmosphere,  and  the  same 
osmotic  pressure  is   shown  by   I  mol  of  any  other  non-electrolyte  in 
the  same  volume.     The  osmotic  pressure  being  proportional  to  the  con- 
centration, if  the  above  sugar  solution  is  concentrated  to  a  volume  of 
I  liter  its  osmotic  pressure  becomes  22.4  atmospheres.     The  osmotic 
pressure  also  increases  with  increasing  temperature  to  the  extent  of  ^i^ 
of  its  amount  at  o°  C.  for  every   i°  rise  in  temperature.     It  follows, 
therefore,  that  the  Laws  of  Boyle  and  Charles  apply  to  osmotic  pressure 
as  well  as  to  gas  pressure.      Not  only  is  this  true,  but  it  is  further 
found  that  the  osmotic  pressure  of  a  given  amount  of  any  non -electrolyte 


Osmotic  Pressure  9 

is  quantitatively  the  same  as  its  gas  pressure  would  be  if  it  existed  alone 
in  the  gaseous  condition  in  the  same  volume  as  that  occupied  by  its 
solution,  and  at  the  same  temperature.  This  is  seen  to  be  true  by 
comparing  the  figures  given  above  (page  7)  for  the  gas  pressure  of 
ammonia  and  the  osmotic  pressure  of  sugar.  Avogadro's  Theory 
applies,  therefore,  to  solutions  as  well  as  to  gases,  and  the  osmotic 
pressure  of  a  solution  depends  solely  on  the  number  of  dissolved  mole- 
cules, and  in  no  wise  upon  their  size  or  chemical  nature. 

Evidently,  then,  the  measurement  of  the  osmotic  pressure  of  a 
solution  may  serve  as  a  method  of  measuring  the  number  of  mols  of 
dissolved  substance  present  in  exactly  the  same  way  as  the  measure- 
ment of  gas  pressure  (or  volume)  may  serve,  by  the  aid  of  Avogadro's 
Theory,  to  determine  the  number  of  mols  of  a  gas  contained  in  a  given 
volume. 

It  has  already  been  noted  that  ail  solutions  of  electrolytes  freeze 
at  abnormally  low  and  boil  at  abnormally  high  temperatures  as  com- 
pared with  solutions  of  non-electrolytes.  Similarly,  it  is  found  that 
the  osmotic  pressure  produced  by  electrolytes  is  much  in  excess  of 
that  occasioned  by  corresponding  amounts  of  non-electrolytes ;  and  it 
is  further  found  that  if  the  number  of  mols  in  such  solutions  is  cal- 
culated from  the  measurements  of  osmotic  pressure,  such  substances 
as  sodium  chloride  yield  somewhat  less  than  twice  as  many  mols,  while 
such  as  calcium  chloride  yield  somewhat  less  than  three  times  as  many 
mols,  as  there  are  chemical  mols  according  to  the  formulas  NaCl  and 
CaCl2. 

In  very  dilute  solutions,  where  the  chemical  molecules  are  completely 
broken  up  into  the  part-molecules  or  ions,  the  osmotic  pressure  is,  of  course, 
twice  and  three  times,  respectively,  that  which  would  be  expected  if  these 
electrolytes  were  not  dissociated.  In  the  case  of  solutions  of  such  concen- 
tration as  would  more  commonly  come  under  consideration,  and  in  which, 
due  to  their  considerable  concentration,  it  is  impossible  for  all  the  molecules 
to  be  dissociated  into  ions,  the  measured  osmotic  pressure  would  indicate 
only  about  1.8  and  2.5  times  as  many  mols,  respectively,  as  the  number  of 
chemical  mols  —  figures  which  are  in  close  accord  with  those  obtained  from 
the  study  of  the  lowering  of  the  freezing  point  or  the  raising  of  the  boiling 
point  of  such  solutions. 

8.  This  dissociation  of  the  electrolyte  is  comparable  with  the  behavior 
of  ammonium  chloride  when  it  is  vaporized.  If  the  density  of  its  vapor  is 
measured,  and  the  molecular  weight  is  calculated  from  this  value  by  the  aid 


io  Electrolytic  Dissociation   Theory 

of  Avogadro's  Theory,  it  is  found  to  be  only  one-half  of  that  corresponding 
to  the  formula  NH4C1.  This  fact  is  easily  explained  if  each  NH4Cl-mole- 
cule  is  dissociated  on  heating  into  two  smaller  individuals,  NH3  and  HCl. 
That  this  actually  occurs  may  be  shown  by  placing  a  lump  of  ammonium 
chloride  in  the  middle  of  a  glass  tube  between  two  plugs  of  asbestos  wool, 
and  heating  the  tube  by  means  of  a  Bunsen  flame.  Of  the  two  dissociation 
products  ammonia  is  the  lighter  and  diffuses  more  rapidly  through  the 
asbestos  plugs,  and  if  moistened  bits  of  red  and  blue  litmus  paper  are 
placed  outside  the  asbestos  the  red  will  at  first  be  turned  blue  at  each  end. 
This  is  due  to  absorption  of  ammonia  by  the  moisture  on  the  litmus  paper 
with  the  formation  of  the  alkaline  ammonium  hydroxide.  If  the  heating  is 
continued  until  the  ammonia  is  mostly  driven  off,  the  litmus  papers  will  then 
turn  red,  due  to  an  excess  of  hydrogen  chloride  gas,  which,  dissolved  in  the 
moisture  on  the  litmus  paper,  forms  a  hydrochloric  acid  solution.  Thus, 
on  vaporization,  ammonium  chloride  is  dissociated  into  two  constituents, 
NH4C1  =  NH3  -f-  HCl,  and  the  apparent  molecular  weight  is  the  mean  of 
that  of  the  two  components  of  the  vapor. 

The  dissociation  of  electrolytes  is,  then,  similar  to  the  dissociation  of 
gases  in  that  it  consists  of  a  splitting  up  of  molecules  into  smaller  indi- 
viduals. A  full  appreciation  of  its  nature  can,  however,  be  obtained  only  in 
connection  with  a  study  of  the  chemical  activity  and  the  electrical  conductivity 
of  electrolytes. 

9.  It   is   evident,   then,   from    the    foregoing   statements   that   the 
phenomena  of  the  lowering  of  the  freezing  point,  raising  of  the  boiling 
point,  and  of  osmotic  pressure,  as  exhibited  by  solutions  of  electrolytes 
and  non-electrolytes,  all  point  to  the  same  conclusion,  namely,  that  the 
molecules    of   electrolytes,  as    ordinarily  expressed   by  their   chemical 
formulas,  are  dissociated  into  a  greater  number  of  smaller  individuals, 
while  the  molecules  of  non-electrolytes  are  not  so  dissociated. 

CHEMICAL  ACTIVITY 

10.  As  a  rule,  chemical  reactions  between  electrolytes  take  place 
with  great   rapidity,   while  reactions  involving   a   non-electrolyte    take 
place  only  slowly.     For  example,  solutions  of  silver  nitrate  and  sodium 
chloride  both  conduct  electricity.     If  they  are  mixed,  there  is  instantly 
formed  a  precipitate  of   silver  chloride,  a  white,  insoluble  substance, 
according  to  the  reaction  AgNO3  +  NaCl  =  NaNO3  +  AgCl. 

A  solution  of  chloroform,  CHC13,  is  a  non-conductor.  If  this  solu- 
tion is  mixed  with  a  solution  of  silver  nitrate  no  precipitate  of  silver 
chloride  is  formed  at  first,  and  it  is  only  after  long  standing  that  one 
slowly  appears.  It  is  reasonable  to  assume  that  there  is  in  all  cases 


Electrolytic  Conduction  1 1 

a  tendency  for  silver  and  chlorine  atoms  to  unite.  Whether  or  not 
this  union  will  take  place  readily  depends  upon  the  state  in  which  these 
atoms  exist.  If  they  are  both  in  the  condition  of  ions,  then  there  is 
no  obstacle  to  prevent  an  immediate  combination ;  but  if,  on  the  other 
hand,  one  or  both  are  already  more  or  less  firmly  bound  as  part  of 
undissociated  molecules,  then,  in  order  for  them  to  combine  with  each 
other,  they  must  first  be  torn  apart  from  the  molecules  in  which  they 
exist.  That  chloroform  is  only  with  difficulty  dissociated  to  yield 
chlorine  ions  is  evident,  both  from  the  fact  that  it  does  not  conduct 
electricity  and  that  it  shows  no  immediate  reaction  with  silver  nitrate. 

ELECTROLYTIC  CONDUCTION 

ii.  Polarity  in  Chemical  Compounds A  conception  which  has 

always  been  associated  with  the  existence  of  chemical  compounds  since 
they  were  first  systematically  studied  is  that  of  polarity.  Chemical 
polarity  in  compounds  is  certainly  closely  related  to  electrical  polarity, 
if,  indeed,  there  is  any  distinction.  For  example,  in  the  compound 
sodium  chloride,  sodium  is  the  positive  constituent,  chlorine  the  nega- 
tive, as  becomes  evident  when  a  current  of  electricity  is  passed  through 
this  electrolyte,  either  when  dissolved  in  water  or  when  in  the  molten 
condition.  If  the  fused  salt  is  used,  its  constituents,  sodium  and  chlo- 
rine, collect  at  the  two  electrodes  respectively,  and  since  sodium  is 
attracted  to  the  negative  electrode,  or  cathode,  it  must  be  the  elec- 
trically positive  constituent  of  the  electrolyte.  As  chlorine,  on  the 
other  hand,  collects  at  the  positive  pole,  ox  .the  anode,  it  must  be  the 
electrically  negative  constituent.  Since  the  two  constituents  accumu- 
late at  opposite  extremes  of  the  body  of  liquid  sodium  chloride,  there 
must  have  been  throughout  the  liquid  a  movement  of  sodium  toward 
the  cathode  and  of  chlorine  toward  the  anode. 

Again,  if  an  electric  current  is  passed  through  an  aqueous  solution 
of  hydrochloric  acid,  hydrogen  is  liberated  at  the  cathode  and  chlorine 
at  the  anode.  Pure  water  does  not  conduct  electricity,  neither  does  dry, 
gaseous  hydrogen  chloride.  But  gaseous  hydrogen  chloride  consists 
only  of  the  undissociated  molecules  HC1,  while  experiments  such  as 
have  been  already  mentioned  show  that  the  molecules  of  hydrochloric 
acid  in  aqueous  solution  are  divided  into  the  smaller  individuals  H  and 


OF  THE 
IIMIX/FRSITY 


12  Electrolytic  Dissociation   Theory 

Cl.  In  view  of  the  fact  that  a  current  is  only  found  to  pass  through 
a  solution  when  accompanied  by  a  movement  in  opposite  directions  of 
particles  which  are  known  to  possess  opposite  polarity,  it  is  plain  that 
it  must  be  these  particles  which  carry  the  current.  Thus,  in  the  solu- 
tion of  hydrochloric  acid  the  movement  of  hydrogen  particles,  which 
possess  positive  electrical  charges,  in  the  direction  towards  the  cathode, 
and  of  chlorine  particles,  which  possess  negative  charges,  in  the  direc- 
tion toward  the  anode,  constitutes  the  actual  movement  of  electricity, 
and  it  is  this  which,  in  fact,  is  the  current. 

12.  Ions The  apparent  anomaly  that  molecules  of  sodium  chlo- 
ride are  dissociated  into  sodium  and  chlorine  atoms,  each  of  which 
exist  separately  as  physical  molecules,  but  without  any  of  the  well- 
known  properties  of  sodium  and  chlorine  becoming  manifest,  may  now 
be  readily  explained.  When  the  original  molecule  is  dissociated  upon 
solution  in  water  its  two  atoms  become  possessed  of  equal  and  oppo- 
site charges  of  electricity,  the  sodium  atom  taking  the  positive  charge 
and  the  chlorine  the  negative.1  These  charged  atoms,  each  having  a 
separate  existence,  are  called  ions.  The  electrical  charges  are  of  great 
magnitude,  and  so  affect  the  character  of  the  ions  as  to  account  for  the 
vast  difference  in  their  properties  as  compared  with  the  atoms  of  the 
electrically  neutral  elements. 

An  ion,  according  to  the  electrolytic  dissociation  theory,  is,  then,  a 
simple  atom,  or  group  of  atoms,  forming  in  itself  a  complete  individual, 
and  possessing  a  charge  of  electricity  ;  and  ions  may  be  produced  by  the 
dissociation  of  larger  electrically  neutral  molecules  (that  is,  of  the  mole- 
cules of  compounds  as  ordinarily  expressed  by  their  chemical  formulas) 
into  smaller  ones  bearing,  respectively,  equivalent  amounts  of  positive 
and  negative  electricity.  In  every  solution  the  aggregate  of  charges 
on  the  positive  ions  which  it  contains  must  be  exactly  equalized  by  the 
sum  of  the  charges  on  the  negative  ions  present,  otherwise  the  solution 
as  a  whole  could  not  be  electrically  neutral. 

The  cause  of  the  phenomenon  of  ionization  is  not  at  all  fully  under- 
stood. In  the  case  of  melted  salts  it  is  probable  that  heat,  by  increasing 

1  Since  equal  quantities  of  opposite  kinds  of  electricity  entirely  neutralize  each  other, 
it  is  immaterial  whether  we  regard  the  original  undissociated  molecule  as  possessing  no 
electrical  charges,  or  as  already  possessing  opposite  charges  which  neutralize  each  other  in 
such  a  manner  that  the  molecule  behaves  as  if  uncharged. 


Ions  1 3 

the  velocity  with  which  the  atoms  of  a  molecule  vibrate,  separates  them 
to  such  an  extent  that  they  no  longer  exist  as  molecules,  but  as  ions, 
which  are  to  a  certain  extent  independent  of  each  other.  In  the  case 
of  solutions,  it  is  apparently  some  specific  power  possessed  by  the  sol- 
vent by  virtue  of  which  it  is  able  to  so  force  itself  in  between  the  parts 
of  the  molecules  as  to  separate  them  into  ions.  Different  solvents  vary 
very  greatly  in  the  extent  to  which  they  possess  this  power.  The 
maximum  among  common  solvents  is  reached  in  the  case  of  water, 
while  with  some  others  it  appears  to  be  entirely  lacking,  as  is  illustrated 
by  the  experiments  on  page  77. 

When  electrodes,  connected  with  a  battery  or  a  dynamo,  are  inserted 
in  a  solution  of  hydrochloric  acid,  the  hydrogen  ions,  by  virtue  of  their 
positive  charges,  are  attracted  to  the  negative  electrode ;  these  ions, 
when  they  touch  the  electrode,  give  up  their  electrical  charges,  neutral- 
izing the  negative  electricity  on  the  electrode,  and  then  become  ordinary 
hydrogen  atoms,  which,  in  turn,  unite  in  pairs  to  form  hydrogen  mole- 
cules, and  thus  hydrogen  gas  escapes  at  the  cathode.  In  an  exactly 
similar  manner  negative  chlorine  ions  are  attracted  to  the  anode, 
where  they  give  up  their  charges  of  electricity.  They  thus  become 
neutral  chlorine  atoms,  which,  combining  in  pairs,  form  the  molecules 
of  the  chlorine  gas  which  escapes. 

13.  As  already  stated,  the  electrical  charges  on  the  positive  and 
negative  ions  of  an  electrolyte  must  just  equalize  each  other.  Figure  I 
represents  a  solution  in  which  ten  of  each  kind  of  ions  are  pictured, 
instead  of  the  almost  countless  number  which  are  in  reality  present  in 
the  solution.  It  is  assumed,  at  the  start,  that  the  charge  on  each  posi- 
tive ion  in  the  upper  row  is  equalized  by  the  charge  on  the  negative 
ion  just  below  it. 


FIGURE  i 


14  Electrolytic  Dissociation   Theory 

If  the  electrodes  -\-  and  —  are  charged,  as  represented  in  the  dia- 
gram, from  a  dynamo  or  battery,  the  negative  ion  farthest  to  the  left  is 
attracted  to  the  positive  electrode,  where  it  parts  with  its  charge,  which 
neutralizes  some  of  the  positive  electricity  on  that  electrode.  This 
would  leave  the  first  positive  ion  at  the  left  unneutralized  by  the  oppo- 
site charge  of  any  other  ion,  and  it  is  at  once  repelled  by  the  positive 
electrode  and  moves  toward  the  right,  while  at  the  same  time  the  second 
negative  ion,  being  attracted  by  it,  moves  toward  it  until  the  two  have 
come  sufficiently  close  together  to  equalize  each  other.  The  second 
positive  ion,  being  now  unbalanced,  attracts  the  third  negative  ion  until 
these  two  equalize  each  other ;  and  so  on,  until  the  ninth  positive  ion 
and  the  tenth  negative  ion,  after  neutralizing  each  other,  leave  the  tenth 
positive  ion  to  be  attracted  to  the  negative  electrode,  there  to  give 
up  its  charge  and  neutralize  a  part  of  the  negative  electricity  on  that 
electrode.  Then  the  solution  is  left  in  the  condition  represented  in 
Figure  2.  At  the  surface  of  each  electrode  is  the  discharged  ion, 
which  is  now  an  electrically  neutral  particle. 


©©©©©©©©©O 

FIGURE  2 


The  further  behavior  of  the  discharged  ions  will  be  spoken  of  in 
Section  16.  It  should  be  noted  that  the  solution  is  now  in  the  same 
condition  as  at  the  start,  but  with  one  less  pair  of  ions,  and  the  process 
just  described  is  ready  to  be  repeated.  In  an  actual  case  of  electroly- 
sis, where  the  ions  are  present  in  enormous  numbers,  this  procedure 
repeats  itself  with  almost  infinite  rapidity. 

14.  Movement  of  the  Ions.  —  That  ions  do  actually  move  through 
a  solution  may  be  most  easily  shown  in  the  case  of  some  electrolyte 
both  ions  of  which  are  colored,  so  that  their  motion  through  the 
liquid  is  made  visible. 


Movement  of  the  Ions 


FIGURE  3 


This  may  be  demonstrated  as  follows :  In  the  lower  bend  of 
a  U-tube  (Figure  3)  is  poured  some  fairly  concentrated  solution  of 
copper  bichromate,  CuCr2O7.1  This  salt  in  aqueous  solution  is  dis- 
sociated into  positive  Cu-ions,  which  are  blue,  and  negative  Cr2O--ions, 
which  are  yellow.  The  resultant  color  is 
green.  On  top  of  this  solution  a  lighter 
solution  of  sodium  chloride  is  so  cautiously 
poured  as  to  avoid  mixing  the  layers.  Plati- 
num electrodes  are  inserted  in  the  two  arms  «-| 1-*-| NaCI 

of  the  tube,  and  a  current  is  passed  through 
the  entire  liquid.  After  about  twenty  min- 
utes there  will  have  appeared  just  above  the  \.  *^- CuCr£07 

green  liquid  a  blue  zone  of  Cu-ions  on  the 
side  toward  the  cathode  and  a  yellow  zone 
on  the  side  toward  the  anode,  both  zones  being  perhaps  2  cm.  high. 
The  blue  zone  is  caused  by  the  movement  of  blue  Cu-ions  upward 
from  the  original  boundary  between  the  two  liquids,  as  well  as  the 
movement  of  the  yellow  Cr2O7-ions  downward,  leaving  only  the  blue 
of  the  Cu-ions.  The  yellow  zone  in  the  other  arm  of  the  tube  also 
appears  above  and  below  the  original  boundary.  It  should  be  noted 
that  after  the  Cu-ions  and  Cr2O7-ions  have  moved  into  the  colorless 
liquid  their  electrical  charges  are  still  balanced,  but  it  is  now  brought 
about  by  the  opposite  charges  on  the  Cl-  and  Na-ions,  respectively, 
which  are  derived  from  the  sodium  chloride. 

The  change  in  position  of  the  boundary  between  the  blue  or  yellow 
and  the  colorless  liquids  in  the  above  experiment  is  not  rapid,2  and  the 
slowness  of  this  motion  is  due  to  the  great  friction  offered  by  the  sol- 
vent to  the  motion  through  it  of  particles  as  minute  as  the  ions.  This 
will  be  understood  if  it  is  recalled  that  while  a  small  fragment  of  chalk 
will  drop  rapidly  to  the  floor  under  the  influence  of  gravity,  yet  if  the 
same  quantity  of  chalk  is  suspended  in  the  air  in  the  form  of  chalk 

1  For  the  details  of  this  experiment  arranged  for  lecture  demonstration,  see  Jour.  Am. 
Chem.  Soc.)  22,  p.  729. 

2  If  the  distance  between  the  electrodes  is  100  cm.,  and  the  difference  of  potential 
between  them  is    100  volts,  the  boundary  lines  would  move  about   2.5  cm.  in   an  hour. 
This  is  about  the  average  rate  of  motion  of  an  ion  through  a  solution  where  the  electrical 
potential  is    I  volt   for  each   centimeter  of   distance   between    the    electrodes.     Hydroxyl 
ions,  however,  move  with  about  three  times,  and  hydrogen  ions  with  about  six  times,  this 
velocity. 


1 6  Electrolytic  Dissociation   Theory 

dust,  as  may  be  brought  about  by  striking  together  two  blackboard 
erasers,  a  long  time  is  required  for  the  chalk  dust  to  settle  to  the  floor, 
although  in  the  latter  case  the  force  of  gravity  is  just  as  great  as  in  the 
former. 

15.  Electrical  Charges  upon  the  Ions.  —  In  view  of  the  state- 
ment just  made  concerning  the  slowness  with  which  the  ions  move 
during  electrolysis,  it  may  appear  remarkable  that  electrolytes  should 
convey  the  current  as  readily  as  they  are  known  to  do.  The  explana- 
tion of  this  fact  is  found  in  the  magnitude  of  the  charge  which  each 
ion  bears.  This  is  so  large  that,  notwithstanding  the  sluggishness  of 
the  motion  of  the  ions,  a  sufficient  quantity  of  electricity  is  brought 
by  the  small  number  of  them  which  do  reach  each  electrode  in  a  given 
interval  of  time  to  constitute  a  considerable  current. 

As  has  already  been  briefly  noted,  a  solution  of  an  electrolyte  gives 
no  external  evidence  of  being  electrified,  such,  for  example,  as  is  shown 
by  a  stick  of  sealing  wax  which  has  been  rubbed  with  a  cat's  skin ; 
that  is,  it  does  not  attract  or  repel  light  bodies  which  are  brought  near 
it.  It  cannot,  therefore,  contain  an  excess  of  either  positive  or  negative 
electricity.  Since,  however,  the  ions  are  assumed  to  carry  charges,  it 
must  follow  that  with  any  electrolyte  the  charges  upon  the  opposite  ions 
must  be  exactly  equivalent  in  amount,  so  that  their  separate  effects  are 
equalized  within  the  solution.  If  the  charge  which  the  hydrogen  ion 
bears  be  taken  as  the  unit  quantity  of  positive  charge,  then  the  chlorine 
ion  must  bear  one  unit  of  negative  charge,  since  hydrochloric  acid  gives 
equal  numbers  of  hydrogen  and  chlorine  ions  :  — 

HCI  -  ->  H+  +  cr. 

From  the  same  reasoning,  it  appears  that  the  sodium  ion  must  bear  one 
unit  of  positive  electricity,  and  the  calcium  ion  two  units  :  — 

NaCl »  Na+    +  Cl~; 

CaCl2  -— >  Ca++  +  Cl~  +  Cl~. 

It  should  be  noted  that,  in  general,  the  number  of  unit  charges  which 
an  ion  bears  is  the  same  as  the  valence  of  the  atom,  or  atom  group, 
from  which  the  ion  is  formed.  A  few  of  the  most  common  ions  with 
the  number  of  units  of  positive  or  negative  charges  which  they  carry 
are  given  in  the  following  table  :  — 


Reactions  at  the  Electrodes  17 

H+  Na+          K+  Ag+  (NH4)+ 

Cl-  Br-  r  (N03)-       (C103r         (CN)-  (C2H3O2)- 

S~-  (S04)"  (Cr207)-- 

Fe+  +  +         A1+  +  +  P(V         (Fe(CN)6)-- 

(Fe(CN)6)-- 

16.  Reactions  at  the   Electrodes In   Section   13    the    process 

has  beer  described  by  which  the  current  is  carried  through  a  solution 
as  a  result  of  the  movement  of  ions  toward  the  electrodes,  where  these 
ions  give  up  their  electric  charges.     The  further  behavior  of  the  dis- 
charged ions  is  something  in  no  wise  connected  with  the  conditions  of 
their  movement  before  they  were  discharged. 

One  of  the  simplest  cases  is  that  of  the  electrolysis  of  a  concen- 
trated hydrochloric  acid  solution.  The  ions  in  such  a  solution,  upon 
being  discharged,  become  simply  electrically  neutral  atoms,  which  at 
this  moment  are  in  the  so-called  nascent  state,  and  are  chemically 
exceedingly  active.  Indeed,  free  uncharged  atoms  of  chlorine  or  hydro- 
gen are  not  capable  of  continued  independent  existence;  if  no  other 
body  is  present  with  which  they  can  unite,  they  themselves  combine  in 
pairs  to  form  the  chemically  much  less  active  molecules  of  chlorine  and 
hydrogen.  These  molecules  then  escape  from  the  liquid  at  the  two 
electrodes  in  the  form  of  chlorine  gas  and  hydrogen  gas. 

17.  Another  typical  case  is  that  of  the  electrolysis  of  a  salt  of  a 
heavy  metal  between  electrodes  of  the  same  metal.     In  a  copper  sul- 
phate solution  the  ions  which  bear  the  current  are  Cu"1"1"  and  SO4~~. 
If   copper  electrodes   are   immersed   in    such  a  solution,  the    positive 
electricity  is  given  up  at  the  cathode  by  the  Cu++-ions,  which  thus 
become  Cu-atoms  and  simply  add  themselves  to  the  mass  of  metal  of 
the  electrode.     The  SO4~  "-ions  about  the  cathode  are  then,  being  elec- 
trically unbalanced,  repelled  by  the  cathode  toward  the  anode,  around 
which  they  collect.     Here,  instead  of  giving  up  their  electric  charges, 
they  find  these  charges  equalized  by  those  on  the  Cu^-ions  which  are 
produced  from  neutral  Cu-atbms  of  the  material  of  the  copper  anode, 
since  it  is  by  the  formation  of  these  positive  ions  out  of  the  mass  of 
the  electrode  that  the  current  enters  the  solution. 


1 8  Electrolytic  Dissociation   Theory 

In  a  case  of  this  kind,  the  fact  that  an  almost  infinitesimal  electrical 
potential  *  is  sufficient  to  cause  some  current  to  pass  through  the  solu- 
tion again  strongly  supports  the  idea  that  electrolytes  are  already  dis- 
sociated into  ions  even  before  a  potential  is  applied.  If  this  were  not 
the  case,  no  current  would  pass  until  a  certain  definite  potential  were 
reached  sufficient  to  pull  the  molecules  apart,  when  the  current  would 
suddenly  begin  to  flow. 

1 8.  It  is  frequently  true  that  the  products  set  free  at  the  electrodes 
give  no  direct  indication  as  to  the  ions   which  conduct  the   current 
through  the  solution.     This  is  illustrated  by  the  electrolysis  of  a  solu- 
tion of   potassium    sulphate   between   platinum    electrodes.     The   ions 
which  transmit  the  current  are  K+K+  and   SO4~~,  but  the  products 
set  free  are  the  gases  hydrogen  and  oxygen.     Furthermore,  the  solu- 
tion at  the  cathode  becomes  alkaline,  while  that  at  the  anode  becomes 
acid.     These  facts  may  be  partially  explained  by  assuming  that  the 
K+-ions  on  reaching  the  cathode  give  up  their  charges  and  become 
electrically  neutral  K-atoms,  which  then  react  with  water  to  produce 
hydrogen  gas  and  potassium  hydroxide ;  the  SO4~  "-ions  on  reaching 
the  anode  are  discharged  and,  not  being  capable  of  existence  alone  in 
the  uncharged  condition,  then  react  with  water  with  the  formation  of 
oxygen  gas  and  sulphuric  acid.     A  more  complete  as  well  as  a  proba- 
bly more  exact  explanation  is  possible,  however,  if  the  Principle  of  Mass 
Action,  together  with  the  fact  that  water  itself  is  slightly  dissociated, 
is  taken  into  consideration.     (See  page  50.) 

19.  A  comprehension  of  the  magnitude  of  the  charges  of  electricity 
which  are  carried  by  the  ions  can  be  obtained  from  the  following  statements 
regarding  the  amount  of  electricity,  the  passage  of  which  through  the  solu- 
tion of  a  metallic  salt,  between  electrodes  of  the  same  metal,  is  accompanied 
by  the  dissolving  from  the  anode  and  depositing  upon  the  cathode  of  an 
equivalent  weight2  in  grams  of  the  metal,  for  example,  31.5  grams  of  'copper. 
This  amount  of  electricity,  if  forced  through  a  50  candle  power  incandescent 
lamp,  would  keep  it  lighted  at  its  full  brilliancy  for  13.5  hours.     If  the  posi- 
tive electricity  discharged  at  the  cathode  by  31.5  grams  of  Cu++-ions  could 
be  condensed  upon  a  metal  sphere,  and  if  the  negative  electricity  produced 
upon  the  anode,  as  a  result  of  the  taking  away  of  the  31.5  grams  of  positive 

1  That  is,  electrical  pressure. 

2  Since  the  copper  ion  carries  a  double  charge  of  electricity,  one-half  an  atomic  weight 
(or  31.5  grams)  of  copper  is  equivalent  to  a  whole  atomic  weight  of  any  element  whose  ion 
carries  but  a  single  charge. 


Summary  19 

Cu~H+-ions  formed,  could  be  condensed  upon  another  metallic  sphere,  and  if 
these  charged  spheres  could  be  held  with  their  charges  upon  them  at  a  dis- 
tance of  3  feet  apart,  the  attraction  pulling  them  together  would  be  equal  to 
a  force  of  io16  tons. 

SUMMARY 

20.  The  conclusions  drawn  from  the  discussions  in  this  chapter 
may  be  summarized  as  follows :  Solutions  have  been  studied  from  five 
standpoints,  namely,  with  reference  to  freezing  point,  to  boiling  point, 
to  osmotic  pressure,  to  chemical  activity,  and  to  electrical  conductivity, 
and  in  each  case  it  appears  that  the  properties  exhibited  by  the  solu- 
tions are  dependent  upon  the  presence  of  the  dissolved  substance  and 
its  condition  within  the  solution.  With  respect  to  this  condition  it  is 
at  once  evident  that  any  substance,  when  dissolved  to  form  a  homo- 
geneous mixture  with  the  solvent,  must  necessarily  be  subdivided  into 
very  small  particles,  which  may,  however,  be  either  larger  or  smaller 
than,  or  identical  with,  the  chemical  molecules.  Of  the  properties 
enumerated,  three,  namely,  freezing  point,  boiling  point,  and  osmotic 
pressure,  are  influenced  only  by  the  number  of  these  particles  present 
in  the  solutions,  and  not  by  their  size  or  chemical  character. 

A  quantitative  study  of  the  three  properties  just  enumerated  has 
shown  that  the  extent  of  the  subdivision  within  the  solution  varies  with 
different  classes  of  substances.  Non-electrolytes,  in  general,  when  dis- 
solved in  water  yield  particles  which  are  identical  with  those  designated 
as  chemical  molecules ;  that  is,  with  those  which  correspond  to  the 
usual  chemical  formulas.  For  example,  ethyl  alcohol,  C2H6O  ;  sugar, 
Ci2H22On  ;  urea,  CON2H4. 

Electrolytes,  on  the  other  hand,  yield  a  number  of  individual 
particles  which  is  greater  than  that  possible  if  only  the  chemical 
molecules  are  considered.  This  leads  directly  to  the  conclusion  that 
some  or  all  of  these  molecules  undergo  further  subdivision  into  smaller 
molecules,  which,  however,  produce  the  same  physical  effect  as  the 
chemical  molecules,  and  to  which  the  name  ions  is  given.  But  it  is 
also  found  that  it  is  only  those  bodies  which,  in  solution,  yield  these 
ions  that  transmit  the  electric  current  and  exhibit  marked  chemical 
activity ;  and  a  consideration  of  these  properties  has  not  only  con- 
firmed the  assumed  independent  existence  within  the  solution  of  these 


20 


Electrolytic  Dissociation  Theory 


parts  of  the  chemical  molecules,  but  it  has  also  indicated  that  each  of 
these  ions  carries  an  electrical  charge  of  very  considerable  magnitude. 
If  the  foregoing  conclusions,  which  are  in  accordance  with  the 
Electrolytic  Dissociation  Theory  as  proposed  by  Arrhenius,  are  well 
established,  it  should  be  found  that  the  values  for  the  degree  of  ioniza- 
tion  of  a  given  substance  in  solution,  determined  from  an  independent 
study  of  such  of  the  different  properties  of  the  solution  as  admit  of 
quantitative  measurement,  should  agree.  The  following  table  shows 
that  these  values  for  certain  well-known  electrolytes  are  concordant, 
the  small  variations  being  satisfactorily  accounted  for  by  what  are 
known  as  unavoidable,  experimental  errors.  In  accordance  with  a  well- 
established  custom,  the  number  of  mols  into  which  one  chemical  mol  is 
dissociated  is  designated  as  i.  The  second  column  of  the  table  shows 
the  molal  concentration  of  the  salt  at  which  the  value  is  determined. 


Salt. 


Molal  conc'n. 


VALUES  OF  /  DETERMINED  FROM 


Osmotic 
pressure. 


Freezing  point. 


Electrical 
conductivity. 


KC1 0.14 

Ca(NO3)2 0.18 

K4(Fe(CN)6) 0.356 

MgS04 0.38 

LiCl 0.13 

SrCl2      0.18 

CaQ2      0.1 

Succinic  acid,  H2C4H4O4   ....  0.1 

Acetic  acid,  HC2H3O2 0.1 

NH4OH     .    . 0.05 


1.81 
2.48 
3.09 
1.25 
1.92 
2.69 
2.47 


1.86 
2.47 

1.20 

1.94 

2.52 

2.65 

1.07 

1.019 

1.038 


1.86 

2.46 

3.07 

1.35 

1.84 

2.51 

2.48 

1.03 

1.013 

1.022 


CHAPTER  II 

THE  LAW  OF  MASS  ACTION  AND  THE  CHEMICAL  BEHAVIOR  OF 

ELECTROLYTES 

REVERSIBLE  REACTIONS  AND  THE  EFFECT  OF  MASS 

21.  In  order  to  understand  the  applications  of  the  Law  of  Mass 
Action,  it  is  first  necessary  to  know  what  is  meant  by  a  reversible 
reaction.  A  familiar  example  is  to  be  found  in  the  action  of  steam 
upon  iron  filings.  If  iron  filings  are  introduced  into  a  tube  and  brought 
to  a  high  temperature,  and  steam  is  then  passed  over  them,  the  iron 
will  be  changed  to  iron  oxide,  and  hydrogen  will  be  liberated  according 

to  the  reaction  3Fe  +  4H2O >  Fe3O4  +  4-H2.  If,  however,  the 

tube  is  kept  at  the  same  temperature  and  hydrogen  is  passed  through 
instead  of  steam,  the  iron  oxide  will  be  reduced  to  iron,  Fe3O4  -+- 
4H2 >  3Fe  +  4H2O;  that  is,  the  first  reaction  is  reversed. 

It  should  be  noted  that  in  the  first  of  these  two  cases  the  steam  is 
supplied  freely,  and  the  hydrogen,  which  is  the  product  of  the  reaction, 
is  forced  out  of  the  tube,  while  in  the  second  case  the  ready  supply 
of  hydrogen  displaces  the  steam  which  is  formed  from  the  reduction  of 
the  iron  oxide. 

Suppose,  however,  that  the  conditions  are  so  altered  that  the  gas- 
eous product  of  the  reaction  cannot  escape,  as  would  be  the  case  if 
the  reactions  were  carried  out  ^n  two  sealed  tubes,  one  of  which  had 
been  previously  charged  with  iron  and  water  vapor,  the  other  with  iron 
oxide  and  hydrogen.  If  these  tubes  were  heated  the  reaction  in  each 
would  begin  as  described  above  for  the  open  tube ;  but  in  the  first  tube 
the  amourft  of  water  vapor  would  gradually  diminish  and  that  of  the 
hydrogen  increase,  while  in  the  second  the  concentration  of  the  hydro- 
gen would  diminish  as  that  of  the  water  vapor  increased.  It  is  easy 
to  see  that  a  point  would  be  reached  in  each  at  which  the  accumu- 
lation of  the  body  which  was  the  product  of  the  initial  reaction  might 

21 


22  Electrolytic  Dissociation   Theory 

be  sufficient  to  cause  a  reversal  of  the  process  to  take  place.  In  fact, 
experiment  has  shown  that,  under  the  conditions  cited,  the  reaction  in 
each  tube  apparently  ceases  at  a  definite  point,  and  that  when  this  point 
has  been  reached  the  ratio  of  the  concentration  of  hydrogen  and  water 
vapor  is  the  same  in  both  tubes.  This  is  called  the  point  of  equilibrium. 
Instead,  however,  of  regarding  this  condition  of  equilibrium  as  one  at 
which  chemical  action  has  ceased,  it  is  better  to  assume  that  both 
reactions  are  still  taking  place,  but  that  by  them  as  many  H2-  as  H2O- 
molecules  are  being  produced  in  a  given  interval  of  time,  so  that  the 
relative  amounts  present  from  moment  to  moment  no  longer  change. 
Such  a  condition  may  be  represented  by  the  equation  3Fe  -f-  4H2O 
^  Fe3O4  +  4-H2,  where  all  four  substances  continue  to  exist  in  equi- 
librium with  each  other. 

22.  From  the  preceding  statements  it  is  evident  that  the  direction 
in  which  a  reversible  reaction  will  proceed  is  determined  by  the  relative 
quantities  of  the  reacting  substances  which  are  present  in  a  given  vol- 
ume, that  is,  upon  their  relative  concentrations.  This  is,  however,  in 
effect  a  statement  of  the  principle  of  the  Law  of  Mass  Action,  which 
may  be  concisely  formulated  as  follows  :  — 

The  rapidity  of  a  c/iemical  cJiange  is  proportional  to  the  concentra- 
tions of  the  substances  taking  part  in  the  reaction. 

Of  the  many  cases  in  which  this  principle  will  find  application  in 
the  remaining  pages,  nearly  all  will  represent  conditions  of  equilibrium 
in  reversible  reactions  where  the  velocities  of  the  opposing  reactions 
are  equal,  as  described  above  for  iron,  steam,  iron  oxide,  and  hydrogen. 

The  law  may  be  restated,  with  special  reference  to  such  conditions, 
as  follows  :  — 

If,  in  a  confined  system,  certain  substances  are  present  which,  by 
reacting  with  one  another,  produce  certain  new  substances,  and  these, 
in  turn,  react  to  reproduce  the  original  substances,  it  is  found  that, 
whatever  the  actual  amounts  of  all  the  reacting  bodies,  the  final  con- 
centrations when  equilibrium  is  reached  always  bear  such  a  relation 
to  each  other  that  the  product  of  the  concentrations  of  the  first  set  of 
substances  stands  in  a  definite  numerical  ratio  to  the  product  of  the 
concentrations  of  the  second  set. 

To  illustrate :    If  the  substances  A  and  B  react  to  form  the  sub- 


Law  of  Mass  Action  23 

stances  C  and  D  according  to  the  chemical  reaction  A  -f-  B  =  C  -f  D, 
into  which  only  one  molecule  of  each  body  enters,  then  whenever  these 
four  bodies  exist  together  in  a  confined  volume  in  a  state  of  equilibrium, 
the  ratio  of  their  concentrations  may  be  expressed  as  follows  :  — 

Cone  n  A  X  Concn  B 

-   — ,         —  =  Constant ', 

Cone  u  C  X  Cone  n  D 

where  Constant  is  a  definite  numerical  quantity  characteristic  of  this 
reaction . 

23.  It  will  be  noticed  that  the  actual  reversible  reaction  considered  in 
Section  21  is  not  so  simple  as  the  ideal  one  just  formulated,  and  it  is  also 
complicated  by  the  fact  that  two  of  the  reacting  bodies,  iron  oxide  and  iron, 
are  solid,  while  the  others  are  gaseous.  The  nature  of  the  experimental  justi- 
fication for  the  Mass  Action  Law  may  be  more  satisfactorily  and,  indeed,  very 
beautifully  demonstrated  by  a  study  of  the  reversible  reaction  involving  the 
dissociation  and  reassociation  of  hydrogen  iodide  at  high  temperatures. 

At  room  temperatures  any  given  volume  of  hydrogen  iodide  gas  may  be 
considered  to  consist  only  of  molecules  corresponding  to  the  symbol  HI. 
If,  however,  a  confined  volume  of  this  gas  is  subjected  to  increasing  tem- 
peratures, it  is  found  that  at  180°  C.,  and  above,  it  undergoes  a  gradual 
dissociation  into  hydrogen  gas  and  iodine  vapor,  according  to  the  equation 
2  HI  =  H.2  -\-  I2,  and  temperatures  may  easily  be  reached  at  which  this 
dissociation  is  practically  complete.  If  now  the  temperature  is  allowed  to 
sink  slowly  there  is  a  gradual  reassociation  of  the  hydrogen  and  iodine 
to  form  hydrogen  iodide ;  and,  moreover,  if  the  composition  of  the  gaseous 
mixture  is  determined  at  any  definite  temperature  during  the  heating  process, 
and  again  at  the  same  temperature  during  the  cooling  process,  the  relative 
amounts  of  iodine,  hydrogen,  and  hydrogen  iodide  are  found  to  be  the  same, 
provided  that  in  either  case  the  temperature  is  held  at  this  point  for  a  suffi- 
cient length  of  time  to  permit  a  state  of  equilibrium  to  become  established. 
It  is  also  true  that  equivalent  amounts  of  hydrogen  and  iodine,  if  heated 
to  a  given  temperature,  will  yield  a  mixture  of  the  three  gases  identical  in 
composition  with  that  obtained  by  starting  with  hydrogen  iodide. 

Experiment  has  shown  that  if  equivalent  amounts  of  hydrogen  and  iodine 
are  heated  to  445°  C.  in  a  sealed  vessel,  by  placing  it  in  the  vapors  of  boiling 
sulphur,  79  per  cent,  of  these  gases  unite  to  form  hydrogen  iodide.  On  the 
other  hand,  if  hydrogen  iodide  is  heated  in  a  similar  manner  approximately 
21  per  cent,  is  decomposed  into  hydrogen  and  iodine. 

The  changes  which  take  place  within  the  gaseous  mixture  may  be 
pictured  somewhat  as  follows :  When  the  mixture  of  hydrogen  and  iodine 
is  suddenly  heated  to  445°  it  contains,  at  the  first  moment,  no  hydrogen 
iodide.  This  body  can  only  form  when  hydrogen  molecules  and  iodine 
molecules  collide  with  each  other,  and  even  then  it  is  necessary  for  them 
to  collide  with  such  a  velocity  and  in  such  a  manner  that  their  atoms  shall 


24  Electrolytic  Dissociation   Theory 

be  so  far  separated  by  the  shock  that,  instead  of  returning  to  their  former 
relative  positions,  they  can  recombine  to  form  different  molecules.  The 
conditions  requisite  for  this  are  established  as  the  temperature  of  the  gas 
increases,  and  a  certain  definite  percentage  of  the  collisions  of  the  different 
molecules  will  result  in  such  an  interchange  of  atoms,  this  percentage  depend- 
ing, in  general,  both  upon  the  temperature  and  upon  the  chemical  nature 
of  the  atoms  and  molecules  concerned.  As  the  number  of  Hi-molecules 
increases  the  number  of  H2-  and  I2-molecules  must  decrease,  consequently 
the  number  of  collisions  between  the  latter  must  also  decrease,  and  the 
amount  of  hydrogen  iodide  which  is  formed  in  a  given  interval  of  time  must 
be  diminished  accordingly.  In  other  words,  the  amount  of  hydrogen  iodide 
formed  in  a  unit  time  (that  is,  the  velocity  of  the  reaction)  depends  upon 
the  number  of  hydrogen  and  iodine  molecules  in  a  unit  volume,  or,  more 
exactly  stated,  it  is  proportional  to  the  concentration  of  H2-molecules  multi- 
plied by  the  concentration  of  I2-molecules.  Representing  the  velocity  of  this 
reaction  by  VT,  it  is  true  that 

Vj  =  Conc'n  H2  X   Conc'n  I2  X   Constant^ 

where  Constanti  is  a  definite  numerical  quantity  depending  on  the  tempera- 
ture and  upon  the  nature  of  this  particular  reaction,  or,  in  other  words,  upon 
the  percentage  of  the  collisions  of  H2-  and  I2-molecules  which  result  in  the 
formation  of  Hi-molecules. 

Again,  when  hydrogen  iodide  is  suddenly  heated  to  445°,  Hi-molecules 
can  only  decompose  when  two  of  them  collide  in  such  a  way  that  both  will 
be  sufficiently  broken  apart  for  a  recombination  of  atoms  into  the  mole- 
cules H2  and  I2  to  take  place.  The  velocity  of  this  reaction,  since,  two 
molecules  of  HI  are  involved,  will  be 

Vn  =  Conc'n  HI  X   Conc'n  HI  X   Constant^ 
or 

Vn  =  (Conc'n  HI)2  X   Constant^. 

As  the  amounts  of  hydrogen  and  iodine  increase,  the  amount  of  hydrogen 
iodide,  and  with  it  the  velocity  of  Reaction  II,  decreases.  Reaction  I  must 
begin  to  take  place  and  constantly  increase  in  velocity  until  21  per  cent,  of 
the  hydrogen  iodide  has  been  converted  into  hydrogen  and  iodine.  At  this 
point  there  is  no  further  change  in  concentrations,  and  the  velocities  of  the 
two  reactions  are  equal,  a  condition  similar  to  that  already  described  in 
the  case  of  the  hydrogen  and  steam.  These  are  the  conditions  of  equilib- 
rium, and  since  under  those  conditions  Vj  and  Vn  are  equal,  the  equivalent 
expressions  must  also  be  equal :  — 

Conc'n  H2  X  Conc'n  I2  X  Constant  =  (Conc'n  HI)2  X  Constant^, 
or 

Conc'n  Ho  X   Conc'n  I2         Constant^ 

—  •=.  —  —  •=.  Co?istant. 

(Conc'n  HI)2  Constant 

This  equation  shows  the  relative  amounts  of  the  different  reacting  substances 
in  a  given  volume  (for  example,  i  liter)  which  must  be  present  when  the 


Degree  of  lonization  25 

point  of  equilibrium   is   reached,  and  is  a  mathematical  expression   of  the 
principle  stated  as  the  Law  of  Mass  Action. 

24.  It  is  probable  that  all  chemical  reactions  are  to  some  extent 
reversible,  and  are  therefore  governed  by  this  law,  which  thus  becomes 
one  of  the  most  important  principles  underlying  chemical  changes.  It 
is  true  that  the  extent  to  which  many  reactions  are  reversible,  under 
conditions  which  usually  prevail,  is  so  small  as  to  be  practically  negli- 
gible ;  but  others,  as  illustrated  above,  may  be  reversed  with  comparative 
ease.  Every  instance  of  the  dissociation  of  an  electrolyte  into  its  ions 
is  such  a  reversible  reaction,  and  the  following  sections  are  devoted  to 
a  discussion  of  the  applications  of  the  Law  of  Mass  Action  to  certain 
familiar  and  typical  instances  of  chemical  change. 

DEGREE  OF  IONIZATION 

25.  It  has  just  been  stated  that  every  instance  of  the  dissociation 
of  an  electrolyte  into  its  ions  is  a  reversible  reaction.  For  example, 
when  sodium  chloride  is  dissolved  in  water  the  condition  of  equilibrium 
which  results  is  represented  by  the  equation  NaCl  ^  Na+  +'C1~. 

From  measurements  of  the  freezing  point  and  of  the  electrical  con- 
ductivity it  is  known  that  in  a  solution  of  sodium  chloride  containing 
o.i  chemical  mol  to  the  liter,  90  per  cent,  of  the  salt  molecules  are 
dissociated  into  ions,  while  10  per  cent,  remain  in  the  undissociated 
condition.  If,  however,  the  volume  of  the  solution  in  which  this  amount 
of  the  salt  is  dissolved  is  varied,  then  the  percentage  of  the  salt  which 
is  dissociated  will  not  remain  the  same,  as  is  evident  from  a  consider- 
ation of  a  mathematical  expression  of  the  Mass  Action  Law  as  applied 
to  this  case,  namely, 

Cone  11  Na+  X  Concn  Cl" 


Concn  NaCl 


=  Constant. 


For,  in  case  the  total  concentration  of  the  salt  were  doubled,  as  may 
be  accomplished  by  evaporating  away  one-half  of  the  water,  then  the 
concentrations  of  all  of  the  three  bodies  Na"1",  Cl~,  and  NaCl  would  be 
doubled,  if  it  were  assumed  that  no  change  in  the  percentage  ionization 
occurred.  Under  these  conditions  the  numerator  of  the  fraction  would 
be  increased  four  times  while  the  denominator  would  be  increased  but 


26  Electrolytic  Dissociation   Theory 

twice,  thereby  doubling  the  value  of  the  fraction  —  a  condition  which 
is  impossible  according  to  the  Law  of  Mass  Action,  since  the  value  of 
the  fraction  must  remain  constant.  In  order  that  it  may  remain  con- 
stant, it  is  necessary  that  the  concentration  of  NaCl  shall  increase  by 
a  greater  amount  than  that  of  either  Na+ or  Cl~;  that  is  to  say,  that 
some  of  the  ions  shall  combine  to  form  undissociated  salt,  and  the 
percentage  of  such  molecules  will,  therefore,  be  greater  than  the  initial 
10  per  cent. 

If  the  total  concentration  of  the  salt  is  diminished,  as  may  be 
accomplished  by  diluting  with  pure  water,  it  is  evident  from  a  similar 
process  of  reasoning  that  more  of  the  undissociated  molecules  must 
dissociate,  and  consequently  the  percentage  of  ionization  must  increase. 

The  principle  just  illustrated  holds  true  for  solutions  of  all  electro- 
lytes, namely,  that  with  increasing  dilution  the  percentage  ionized 
increases  (although,  of  course,  the  actual  concentration  of  the  ions 
must  decrease),  while  with  increasing  concentration  the  percentage 
ionization  decreases. 

THE  FORMATION  OF  INSOLUBLE  COMPOUNDS  THROUGH  THE 
INTERACTION  OF  CERTAIN  IONS 

26.  When  solutions  of  two  or  more  electrolytes  are  mixed  the  ions 
which  they  contain  at  once  enter  into  new  combinations,  with  the 
formation  of  larger  or  smaller  amounts  of  undissociated  molecules  not 
found  in  either  of  the  original  solutions.  Whether  or  not  this  rear- 
rangement of  the  constituents  of  the  solution  is  at  once  made  evident 
to  the  observer  will  usually  depend  upon  the  solubility  of  the  new 
compounds  which  may  have  resulted.  For  example,  in  separate  solu- 
tions of  sodium  chloride  and  potassium  nitrate  the  ions  of  each  salt 
are  in  equilibrium  with  the  undissociated  molecules  of  the  respective 
compounds,  as  expressed  by  the  equations  :  — 

NaCl     ^  Na+  +  Cl~; 
KNO3  ^  K+    +'NO8- 

O     x  I  O 

If  now  these  two  solutions  are  mixed,  a  certain  number  of  undis- 
sociated molecules  of  two  new  salts  will  be  formed,  and  this  inter- 
change will  continue  until  new  conditions  of  equilibrium  are  established 
as  expressed  by  the  equations  :  — 


Interaction  of  Ions  .  27 

K+    +      Cr^KCl; 
Na+  +  NO3~  F^  NaNO3. 

In  this  particular  instance  the  observer  will  discover  no  evidence 

that  any  such  change  has  occurred,  and,  indeed,  the  actual  change  is 
I  of  little  magnitude,  because,  first,  the  new  compounds,  potassium  chlo- 
[  ride  and  sodium  nitrate,  as  well  as  the  original  salts,  sodium  chloride 

and  potassium  nitrate,  are  strong  electrolytes,  and  hence  largely  ionized ; 

and,  second,  because  all  are  very  soluble  bodies,  and  therefore  remain 

in  solution. 

27.  On  the  other  hand,  if  any  new  combination  of  ions  can  result 
in  the  production  of  an  insoluble  body,  the  resulting  changes  may  be  of 
much  greater  magnitude.     For  example,  in  solutions  of  potassium  chlo- 
ride and  silver  nitrate,  taken  separately,  a  condition  of  equilibrium  exists 
between   their   undissociated  molecules  and  ions,  in  which,  as   in  the 
cases  cited  in  the  previous  paragraph,  the  ions  very  largely  predominate. 
If,  however,  these  solutions  are  mixed,  a  combination  both  of  Ag+-ions 
with  Cl~-ions  to  form  undissociated  silver  chloride,  and  of  K+-ions  with 
NO3~-ions  to  form  undissociated  potassium  nitrate,  becomes  possible ; 
and,  since  the   silver  chloride  is  almost  completely  insoluble,  the  mo-' 
ment  that  any  of  this  compound  is  formed  it  separates  from  solution 
in  the  form  of  a  precipitate.      In  this  way  precipitation  will  continue 
till  either  the  Ag+-  or  Cl~-ions,  or  both,  are  used  up,  because  silver 
chloride,  being  insoluble,  cannot  remain  in  solution  in  the  undissociated 
form  in  sufficient  quantity  to  maintain  equilibrium  with  any  appreciable 
quantity  of  its  ions. 

28.  The  remaining  solution  contains  K+-  and  NO3~-ions,  with  but 
a   very    small   proportion   of   undissociated    KNO3.       If    this    solution 
is  filtered   from    the   precipitate   and    evaporated   until  an   insufficient 
quantity  of  water  is  left  to  keep  the  undissociated  potassium  nitrate  in 
solution,  this  salt  then  separates  as  crystals  from  the  solution.     It  is  to 
be  noted  that  this  separation  is  entirely  comparable  with  the  precipita- 
tion of  the  silver  chloride  just  described,  the  difference  in  their  manner 
of  formation  being  wholly  due  to  the  difference  in  their  solubilities. 

It  may  be  stated  as  a  general  rule,  which  has  but  very  few  excep- 
tions, that  salts  in  solution  are  largely  dissociated  into  their  ions.  This 
being  true,  it  is  evident  that  there  will  be  but  slight  chemical  change 


28  Electrolytic  Dissociation   Theory 

whenever  solutions  of  salts  are  mixed,  unless  one  of  the  new  salts  which 
may  be  formed  by  the  mixture  is  very  little  soluble. 

SOLUBILITY  PRODUCT 

29.  When  any  pure,  solid  substance  is  brought  into  contact  with  a 
liquid  in  which  it  is  soluble,  its  molecules  immediately  begin  to  pass 
from  its  surface  into  the  liquid,  much  as  the  molecules  of  a  gas  pass 
into  a  vacuous  space,  or  the  molecules  of  a  liquid,  as  water,  pass  into 
the  atmosphere  when  it  evaporates.  This  process  continues,  if  the  solid 
is  present  in  sufficient  quantity,  until  a  solution  results  which  is  satu- 
rated with  the  solid  at  the  temperature  employed.  Under  the  conditions 
then  prevailing  the  molecules  of  the  solid  still  continue  to  pass  into  the 
solvent  liquid,  but  molecules  are  also  returning  from  the  solution  to 
the  solid,  and  the  number  passing  in  each  direction  has  become  exactly 
equal.  Consequently  there  are  no  further  changes  in  concentration  of 
the  solution,  and  a  condition  of  equilibrium  exists  between  the  solid 
and  its  dissolved  molecules.  The  amount  of  any  substance  which  has 
dissolved  to  form  a  saturated  solution  is  always  a  definite  quantity  for 
a  given  temperature,  and  is  characteristic  of  that  particular  substance. 
It  may  be  expressed  as  the  number  of  grams  which  are  dissolved  in 
a  liter  of  solution. 

A  non-electrolyte,  such  as  sugar,  dissolves  in  water  solely  in  the 
form  of  undissociated  molecules  until  these  have  reached  the  saturation 
concentration,  when  they  are  in  equilibrium  with  solid  sugar.  An  elec- 
trolyte also  dissolves  as  undissociated  molecules  at  first,  and  when  the 
solution  is  saturated  these  undissociated  molecules  will  be  in  equilibrium 
with  the  solid  salt,  and  will  have  a  certain  definite  concentration,  just  as 
the  molecules  of  undissociated  sugar  had  a  certain  definite  concentra- 
tion when  in  equilibrium  with  the  solid  sugar.  But  the  case  of  the 
electrolyte  is  further  complicated  because  its  molecules,  as,  for  example, 
those  of  silver  bromate,  as  soon  as  they  dissolve  dissociate  for  the  most 
part  into  ions,  whereby  the  concentration  of  the  un-ionized  molecules 
would  be  diminished  were  it  not  for  the  fact  that  more  solid  salt  at  once 
dissolves  until  the  equilibrium  is  reestablished.  When,  therefore,  solid 
silver  bromate  is  suspended  in  water  it  will  dissolve  until  that  concen- 
tration of  the  silver  bromate  molecules  is  reached  which  stands  in 


Solubility  Product  29 

equilibrium  with  the  solid,  and  also  until  that  concentration  of  the 
silver  and  bromate  ions  is  reached  which  stands  in  equilibrium  with 
the  undissociated  molecules.  This  is  expressed  by  the  equation:  — 

AgBr03  ^  AgBr03  ^  Ag+  +  BrO3-. 

The  relation  between  the  concentrations  of  the  ions  and  molecules 
in  the  solution  may  be  expressed  according  to  the  Mass  Action  Law  as 
follows  :  — 

Cone 'n  Ag+  X  Concn  BrO3~ 

— — ; —  -  —  Constantly 

Cone  n  AgBrO3 

or 

Concn  Ag+  X  Concn  BrO3~  =  Constant -^  X  Concn  AgBrO3. 

Since  in  a  saturated'  solution  the  concentration  of  undissociated  silver 
bromate  is  always  the  same,  the  value  of  Concn  AgBrO3  is  also 
a  constant  quantity,  and  the  relation  expressed  above  may  be  simplified 
to 

Cone  n  Ag+  X  Concn  BrO3~  =  Constant ; 

that  is  to  say,  in  a  saturated  solution  of  silver  bromate  the  concentra- 
tion of  the  silver  ions  multiplied  by  the  concentration  of  the  bromate 
ions  is  equal  to  a  definite  quantity.  This  quantity  is  called  the  solu- 
bility product  for  silver  bromate.  The  general  application  of  this 
principle  may  be  stated  as  follows :  In  a  saturated  solution  of  any 
electrolyte  the  product  of  the  concentrations  of  its  ions  has  a  certain 
definite  value,  which  is  known  as  the  solubility  product  for  that  sub- 
stance.1 If  this  value  is  exceeded  the  solution  is  supersaturated,  and 
some  of  the  solid  substance  must  separate  from  the  solution.  If  this 
value  has  not  been  reached  more  of  the  solid  substance  can  dissolve. 
30.  The  wider  significance  of  the  solubility  product  becomes  evi- 
dent from  the  behavior  of  a  saturated  salt  solution  when  other  salts  are 
dissolved  in  it.  Thus,  if  some  crystals  of  a  soluble  salt,  such  as  sodium 
nitrate,  which  gives  no  ion  in  common  with  silver  bromate,  are  dropped 
into  a  solution  of  that  salt  and  stirred  about,  there  will  be  no  change 

1  In  the  case  of  an  electrolyte,  the  molecule  of  which  yields  more  than  one  ion  of 
a  given  sort,  the  concentration  of  that  ion  must  be  raised  to  a  higher  power,  e.g., 
PbI2  ^±  Pb++  +  I-  +  I~,and  the  solubility  product  becomes  Conc'n  Pb++  X  (ConSn  I~)2. 


30  Electrolytic  Dissociation   Theory 

noticed  except  that  the  added  salt  dissolves  with  the  same  readiness  as 
in  pure  water.  If,  on  the  other  hand,  crystals  of  silver  nitrate,  a  salt 
which  does  have  an  ion  in  common,  are  used,  they  dissolve  as  in  pure 
water,  but  immediately  following  their  solution  there  separates  out  all 
through  the  solution  a  finely  crystalline  precipitate  of  silver  bromate. 
Crystals  of  potassium  bromate,  when  dissolved,  also  produce  an  exactly 
similar  precipitation  of  silver  bromate.  This  may  be  explained  as  fol- 
lows :  In  the  saturated  solution  containing  only  silver  bromate  the 
concentration a  of  both  the  Ag+-  and  BrO3~-ions  is  the  same,  and  their 
product  is  the  solubility  product.  The  value  of  this  product  is  in  no 
wise  altered  by  the  presence  in  the  solution  of  moderate  amounts  of 
other  indifferent  ions.  But  when,  with  the  addition  of  silver  nitrate, 
an  excess  of  Ag+-ions  is  brought  into  the  solution,  the  product  of  the 
concentrations  of  the  ions  is  increased  beyond  the  solubility  product ; 
that  is,  the  solution  is  supersaturated,  and  the  product  can  only  be 
brought  back  to  its  normal  value  by  the  disappearance  of  Ag+-  and 
BrO3~-ions  together,  as  they  unite  to  form  the  undissociated  salt. 
Since,  because  of  its  slight  solubility  in  water,  only  a  very  small  amount 
of  the  latter  can  exist  in  solution,  it  continues  to  separate  in  the  solid 
state  as  fast  as  it  is  formed.  If,  for  example,  sufficient  silver  nitrate 
has  been  added  to  make  the  final  concentration  of  Ag+-ions  ten  times 
as  great  as  at  the  start,  then  only  one-tenth  of  the  original  BrO3~-ions 
can  remain  in  the  solution.  If  instead  of  Ag+-ions  an  excess  of  BrO3~- 
ions  is  brought  into  the  solution  by  the  addition  of  potassium  bromate, 
the  concentration  of  the  Ag+-ions  must  be  diminished  through  separation 
of  the  solid  salt  in  an  exactly  similar  way. 

Likewise  when  solid  silver  bromate  is  stirred  into  pure  water,  or 
solutions  of  silver  nitrate  or  potassium  bromate,  it  will  dissolve  only 
until  the  solubility  product  is  reached,  and  the  amount  which  can  dis- 
solve before  this  product  is  attained  is  naturally  less  in  the  last  two 
cases  than  in  the  first,  because  of  the  previous  existence  within  these 
solutions  of  Ag+-  or  BrO3~-ions.  The  few  figures  which  follow,  taken 
from  actual  experiments,  are  in  entire  agreement  with  the  theory:  — 

1  By  concentration,  in  such  a  case  as  this,  is  meant  the  number  of  equivalent  weights  in 
grams  of  the  ions  or  salts  in  question,  rather  than  the  number  of  grams  which  are  contained 
in  a  definite  volume  of  the  solution.  For  example,  in  a  solution  containing  108  grams  of 
Ag+-ions  and  128  grams  of  BrOs~-ions,  the  concentration  of  the  two  ions  would  be  said  to 
be  the  same,  as  these  quantities  represent  an  equivalent  weight  in  each  case. 


Reactions  of  the  Ions  31 

Solubility  of  AgBrO3  in  pure  water  =  0.00810  molal.1 

Solubility  of  AgBrO3  in  0.0364  molal  KBrO3    =  0.00227  molal. 
Solubility  of  AgBrO3  in  0.0364  molal  AgNO3  =  0.00216  molal. 


CHARACTERISTIC  REACTIONS  OF  THE  VARIOUS  IONS 

31.  Simple  Ions It  has  already  been  stated,  on  page  27,  that 

when  a  solution  containing  Ag+-ions,  such  as  a  solution  of  silver  nitrate, 
is  mixed  with  one  containing  Cl~-ions,  such  as  potassium  chloride,  these 
ions  unite  at  once  to  form  the  characteristic  compound,  silver  chloride, 
which,  because  of  its  insolubility  in  water,  is  thrown  out  of  solution  as  a 
precipitate.     Since  this  reaction  always  occurs  when  Ag+-  and  Cl~-ions 
are  brought  together,  it  may  be  used  to  demonstrate  the  presence  of 
either  of  these  ions  in  a  solution  ;  and  all  substances,  such  as   HC1, 
NaCl,  MgCl2,  A1C13,  etc.,  which  yield  Cl~-ions  should  cause  this  reac- 
tion.    For  example,  if  the  addition  of  a  solution   of  silver  nitrate  to 
another  solution  occasions  the  precipitation  of  a  white,  curdy  substance, 
which  turns  dark  on  exposure  to  a  strong  light,  then  this  substance  is 
probably  silver  chloride  and  the  solution  contains  Cl~-ions. 

It  is,  of  course,  equally  true  that  this  reaction  may  be  used  as  a 
test  for  Ag+-ions  by  adding  a  solution  of  HC1,  for  example,  to  the 
solution  suspected  to  contain  the  Ag+-ions. 

Moreover,  it  is  possible  to  select  reactions  which  are  similarly  char- 
acteristic of  a  large  number  of  the  ions ;  thus  Ba+  +-ions  when  brought 
into  contact  with  SO4~~-ions  always  yield  a  white  precipitate  of  barium 
sulphate;  Pb++-ions  when  brought  into  contact  with  S~~-ions  cause  a 
black  precipitate  of  lead  sulphide  to  fall,  and  so  on.  These,  as  will  be 
noted  later,  constitute  the  reactions  upon  which  the  systematic  scheme 
of  qualitative  chemical  analysis  is  based,  as  well  as  the  procedures  of 
quantitative  analysis. 

32.  In  the  illustration  of  the  application  of  the  principle  of  the  solubility 
product  cited  in  Section  30,  the  salts  which  produced  the  increased  concen- 
tration of  the  Ag+-  and  BrO8~-ions  were  assumed  to  be  added  to  the  solution 
in  the  solid  form.     A  little  reflection  will  show  that  if  such  salts  are  readily 
soluble  in  water,  thus  permitting  the  preparation  of  solutions  which  are  highly 

1  The  solubility  is  here  expressed  as  the  number  of  chemical  mols  per  liter  (see  also 
the  previous  footnote). 


32  Electrolytic  Dissociation   Theory 

concentrated  as  compared  with  the  solution  of  the  slightly  soluble  silver 
bromate,  the  addition  of  these  concentrated  solutions  would  tend  to  produce 
a  similar  effect  to  that  resulting  from  the  addition  of  the  solid  salt,  namely, 
to  lessen  the  solubility  of  the  bromate.  The  reagents  used  in  qualitative  and 
quantitative  analysis  are  usually  just  such  relatively  concentrated  solutions  of 
various  salts,  and  the  substances  which  they  precipitate  are  bodies  of  even 
less  solubility  than  the  silver  bromate.  It  is  easy,  then,  to  understand  that, 
in  general,  the  test  for  such  an  ion  as  SO4~  ~  (or  the  quantitative  measure- 
ment of  that  ion)  by  the  addition  of  Ba++-ions  (precipitation  of  BaSO4)  may 
be  made  more  delicate  by  the  addition  of  a  moderate  excess  of  the  reagent ; 
that  is,  by  increasing  the  concentration  of  the  Ba++-ions,  thus  rendering  the 
separation  of  the  BaSO4  more  complete. 

33.  Complex  Ions Not  all  substances  which  contain  chlorine  or 

silver  will,  however,  yield  Ag+-  or  Cl~-ions  if  tested  as  above  described. 
The  failure  may  be  due,  as  in  the  case  of  chloroform,  to  the  fact  that 
no  dissociation  occurs,  or  it  may  be  due  to  the  formation  within  the 
solution  of  other  than  simple  Ag+-  and  Cl~-ions.  For  example,  the  salt 
silver  ammonia  nitrate  can  be  obtained  by  adding  an  excess  of  ammonia 
water  to  a  silver  nitrate  solution,  Ag+,  NO3~  +  2NH3  =  (Ag.2NH3)+, 
NO3~.  This  salt  is  undoubtedly  ionized,  but  in  such  a  way  as  to  yield  a 
complex  ion  which  exhibits  properties  very  different  from  those  of  simple 
Ag+-ions,  and  does  not  combine  with  Cl~-ions  to  form  an  insoluble  salt. 

In  a  similar  way  complex  anions  may  be  formed,  and  some  elements 
which,  as  simple  ions,  behave  as  cations  may  become  a  part  of  such 
anions.  The  salt  potassium  silver  cyanide  will  serve  to  illustrate  this. 

As  silver  cyanide  is  insoluble,  Ag+-ions  and  CN~-ions  when  brought 
into  the  same  solution  cause  a  precipitate  of  AgCN  to  form.  If,  how- 
ever, this  precipitate  is  treated  with  an  excess  of  potassium  cyanide 
solution,  the  molecules  of  AgCN  have  the  power  of  combining  with 
CN~-ions  in  the  solution  to  form  a  new  complex  ion.  This,  as  well 
as  the  K+-ions,  is  soluble,  and  the  precipitate  redissolves,  K+,  CN~  + 
AgCN  =  K+,  (Ag(CN)2)-. 

That  the  silver  is  a  constituent  of  the  positive  ion  of  silver  ammonia 
nitrate  and  of  the  negative  ion  of  the  potassium  silver  cyanide  is  shown 
by  the  fact  that  when  an  electric  current  is  passed  through  solutions  of 
the  two  salts,  the  ion  containing  the  silver  moves  in  the  first  case  toward 
the  cathode,  and  in  the  second  toward  the  anode.1 

1  The  details  of  an  experiment  to  illustrate  this  may  be  found  in  Jour.  Am.  Chem.  Soc., 
22,  p.  732. 


Acids  33 

All  of  the  elements  which  are  capable  of  forming  chemical  com- 
pounds, that  is,  all  except  those  of  the  argon  group,  can  exist  in  some 
sort  of  ionic  condition.  Many  yield  simple  ions,  with  one  or  more 
electrical  charges,  while  others  can  only  form  a  part  of  complex  ions. 
For  example,  carbon  and  nitrogen  cannot  form,  in  any  appreciable 
concentration,  ions  consisting  of  single  charged  atoms,  but  they  can 
form  with  great  readiness  such  ions  as  CO3~~,  NH4+,  NO2~,  and  NO3~. 

A  list  of  the  ions  of  the  more  common  elements  which  are  ordinarily 
met  with  in  the  course  of  qualitative  analysis,  together  with  their  most 
important  characteristics,  will  be  found  in  Chapter  V. 

34.  It  should  be  pointed  out  here  that  complex  ions,  both  anions  and 
cations,  are  formed  in  great  variety,  and  that  their  formation  frequently  in- 
creases the  apparent  difficulty  in  the  application  of  the  Law  of  Mass  Action 
to  specific  instances  of  chemical  change.     These  difficulties  lessen,  however, 
with   the   increase    of   our  knowledge;   and    when  the  exact  conditions  are 
unknown  in  a  special  case,  it  is  frequently  possible  to  draw  fairly  accurate 
conclusions  from  analogies. 

35.  Acids All  acid  solutions  possess  certain  characteristic  prop- 
erties :  (i)  They  turn  blue  litmus  red;  (2)  they  taste  sour;    (3)  they 
dissolve  certain  metals  with  an  evolution  of  hydrogen  gas ;  (4)  they  neu- 
tralize bases ;  (5)  they  conduct  electricity.     All  of  these  properties  of 
solutions  of  acids  are  those  of  one  common  component,  the  hydrogen 
ion,  and  an  acid   may  be  defined  as  a  substance  which,  in  aqueous 
solution,  yields  hydrogen  ions. 

It  is  well  known  that  acids  vary  in  strength,  that  is,  some  yield 
solutions  which  are  more  intensely  sour  than  others,  or  react  more 
vigorously  with  such  a  metal  as  zinc  or  magnesium  ;  and  a  study  of 
the  solutions  of  the  various  acids  has  shown  that  these  variations  in 
strength  are  due  to  variations  in  dissociation.  The  strongest  (most 
active)  acids  are  those  which  are  most  dissociated,  that  is,  those  which 
yield  the  largest  relative  number  of  H+-ions ;  and,  since  it  has  already 
been  stated  that  the  electrical  conductivity  is  also  dependent  upon  the 
degree  of  dissociation,  it  is  evident  that  the  strength  of  an  acid  may  be 
conveniently  determined  by  measuring  its  conductivity. 

36.  This  may  be  approximately  accomplished  with  the  aid  of  an  appa- 
ratus designed  by  Dr.  W.  R.  Whitney  and  illustrated  in  the  accompanying 
figure.1     In  the  four  vertical  tubes  are  to  be  placed  four  acids  of  different 

1  For  a  full  description  of  the  details  of  this  experiment  see  Jour.  Am.  Chem.  Soc.,  22, 
P-  736. 


34 


Electrolytic  Dissociation   Theory 


FIGURE  4 


degrees  of  ionization.  In  each 
tube  are  two  electrodes,  the  up- 
per movable,  the  lower  rigid ; 
and  in  series  with  each  tube  is 
an  ordinary  incandescent  lamp. 
Connections  are  made  with  a 
lighting  circuit  (alternating  cur- 
rent, ii2-volt  is  best),  so  that  a 
current  may  pass  through  each 
tube,  the  magnitude  of  the  cur- 
rent being  judged  by  the  bril- 
liancy with  which  the  lamp  below 
that  tube  glows.  If  ^  normal 
solutions  of  hydrochloric  acid, 
sulphuric  acid,  chloracetic  acid, 
and  acetic  acid  are  placed  in  the 
four  tubes,  and  the  electrodes  are 
all  placed  at  the  same  distance 
apart,  the  lamp  in  series  with  the 
hydrochloric  acid  glows  bril- 
liantly, that  with  sulphuric  acid 
a  little  less  brilliantly,  that  with 


chloracetic  acid  much  less  brilliantly,  while  that  with  acetic  acid  barely  glows. 
If  now  the  upper  electrodes  are  adjusted  so  that  all  the  lamps  glow  with 
equal  brightness,  it  will  be  found  that  the  distances  between  the  two  elec- 
trodes bear  approximately  the  ratios  TOO  :  85  :  15  :  i  ;  that  is,  the  relations 
of  these  four  acids,  with  respect  to  conductivity,  degree  of  dissociation,  and 
strength,  are  approximately  those  indicated  by  these  figures. 

37.  Bases.  —  All  basic  solutions  have  the  following  characteristic 
properties:    (i)  They  turn  red    litmus  blue;    (2)  they  taste  alkaline; 
(3)  they  cause  a  slippery  feeling  between  the  ringers ;    (4)  they  neu- 
tralize acids ;  (5)  they  conduct  electricity.     These  are  all  properties  of 
one  common  component,  the  OH~-  or  hydroxyl-ion.     The  strength  of  a 
base  depends  upon  the  extent  to  which  it  is  dissociated ;  that  is,  upon 
the  relative  number  of  OH  "-ions  which  it  yields. 

With  the  same  apparatus  used  above,  the  relative  strength  of  potassium 
hydroxide,  sodium  hydroxide,  and  ammonium  hydroxide  may  be  shown  to 
be  approximately  100  :  100  :  i. 

38.  Salts If  separate  solutions  of  an  acid  and  a  base  are  mixed 

together  there  is  an  immediate  reaction,  and  if  the  amounts  are  properly 
chosen  the  resulting  solution  is  exactly  neutral ;  that  is,  it  shows  none 
of  the  characteristic  properties  of  the  H+-  or  the  OH  "-ions.     The  expla- 
nation of  this  becomes  clear  from  an  inspection  of  a  reaction  represent- 


Neutralization  35 

ing  such  a  process  of  neutralization,  Na+,  OH"  +  H+,  Cl~  =  HOH  + 
Na+,  Cl~,  which  shows  that  the  principal  and  important  change  in  the 
constituents  of  the  solution  has  resulted  in  the  combination  of  the  H+- 
and  OH~-ions  to  form  water.  As  water  is  practically  undissociated, 
these  ions  are  no  longer  in  a  condition  to  impart  their  characteristics  to 
the  solution.  The  remaining  constituents  of  the  solution  are  Na+-  and 
Cl~-ions,  and  if  this  solution  is  evaporated  the  dissociation  of  the  NaCl 
gradually  lessens  as  the  solution  becomes  more  concentrated,  and  finally 
the  solid  salt  separates  exactly  as  was  described  in  the  case  of  the 
potassium  nitrate,  on  page  27. 

A  salt  is  thus  a  product  of  the  neutralization  of  an  acid  and  a  base, 
and  is  formed  from  the  negative  ion  of  the  former  and  the  positive  ion 
of  the  latter.  As  has  already  been  stated,  salts  in  solution  (with  very 
few  exceptions)  are  highly  ionized. 

If,  after  showing  the  relative  degree  of  ionization  of  the  four  acids  as 
above  described,  the  upper  electrodes  are  removed,  and  after  adding  a  drop 
of  litmus,  a  solution  of  potassium  hydroxide  is  added  drop  by  drop  to  each 
tube  until  the  color  barely  changes  to  blue,  it  will  be  found,  on  reinserting 
the  upper  electrodes  to  points  one-third  of  the  distance  from  the  bottom  of 
each  tube,  that  all  the  lamps  glow  brilliantly  and  with  approximately  the 
same  intensity.  This  shows  that  the  four  salts  produced  by  the  neutraliza- 
tions, namely,  potassium  chloride,  potassium  sulphate,  potassium  chlorace- 
tate,  and  potassium  acetate,  are  all  ionized  to  approximately  the  same 
extent,  regardless  of  the  acid  from  which  they  are  formed. 

39.  Neutralization The  essential  reaction  in  the  neutralization 

of  an  acid  and  a  base  is  the  mutual  disappearance  of  H+-  and  OH~- 
ions.  Where  both  acid  and  base  are  highly  ionized  and  the  resulting 
salt  is  soluble,  this  is,  as  already  stated,  practically  the  only  reaction. 
The  truth  of  this  statement  is  indicated  by  the  fact  that  the  heat 
evolved  by  the  neutralization  of  equivalent  quantities  of  the  following 
strong  acids  and  bases  is  in  all  cases  the  same.1 


HC1 

HNO3 

NaOH 

13,700 

13,700 

KOH    .......... 

13,700 

13,800 

1  A  calorie  is  the  amount  of  heat  which  is  sufficient  to  raise  the  temperature  of  I  gram 
of  water  through  i°  C.  The  figures  given  in  the  table  are  the  number  of  calories  which 
are  produced  by  the  neutralization  of  molecular  amounts  of  the  various  acids  and  bases. 


36  Electrolytic  Dissociation   Theory 

The  heat  of  the  reaction  H+  -f-  OH"  =  H2O  is,  therefore,  13,700 
calories. 

As  would  be  expected,  the  mixing  of  neutral  salt  solutions,  as 
potassium  nitrate  and  sodium  chloride,  where  the  ionic  changes  are 
exceedingly  small,  produces  no  appreciable  heat  effect. 

40.  When,  however,  either  the  acid  or  base  is  weak  we  find  that 
the  heat  of  neutralization  is  quite  different :  — 


HC1 

HN03 

H(C2H;J02) 

H2S 

NaOH    

13,300 

3,800 

KOH 

13  300 

3800 

NH4OH         .... 

12,400 

12500 

1L  000 

3  100 

From  what  has  been  said  above  it  is  evident  that  the  lessening 
of  the  amount  of  heat  which  is  evolved  in  these  cases  must  be  due 
to  some  ionic  changes  within  the  solution  which  themselves  consume 
energy.  The  nature  of  these  changes  can  best  be  learned  from  a  study 
of  a  typical  case,  that  of  the  neutralization  of  acetic  acid  and  ammonium 
hydroxide. 

The  dissociation  of  both  of  these  electrolytes  in  moderately  dilute 
solutions  does  not  exceed  I  per  cent.  The  conditions  of  equilibrium 
within  the  separate  solutions  are  represented  by  the  expressions  :  — • 

Concn  H+  X  Concn  Ac" 


Cone  n  HAc 
Conc'n  NH4+  X  Conc'n  OH 


=  Constant  ; 
=  Constant. 


Concn  NH4OH 

(The  values  of  the  two  constants  are  not,  of  course,  the  same.     The 
first  is  found  to  be  0.000018;  the  second,  0.000023.) 

If  now  such  quantities  of  the  two  solutions  are  mixed  as  will 
exactly  neutralize  each  other,  a  series  of  changes  occurs.  At  first  the 
H+-  and  OH  "-ions  unite  to  form  undissociated  water  exactly  as  in 
the  case  of  strong  electrolytes,  and  it  requires  but  a  moment's  reflec- 
tion to  see  that  this  alters  the  concentration  of  these  ions  within  the 


Neutralization  37 

solution,  and  that,  in  order  that  the  equilibrium  condition  may  be 
maintained,  more  H+-  and  OH~-ions  must  be  furnished  ;  that  is,  that 
the  undissociated  HAc-  and  NH4OH-molecules  must  dissociate:  — 

HAc  ^  H+  +  Ac"; 
NH4OH^NH4+-f  Oil- 
But  the  H+-  and  OH~-ions  thus  formed  also  combine,  and  this  pro- 
cedure is  repeated  with  great  velocity  until  the  solution  is  neutral,  and 
complete  equilibrium  between  all  of  the  components  is  established ; 
that  is,  until  all  of  the  HAc-  and  NH4OH-  molecules  have  been  dis- 
sociated. At  this  point  the  solution  will  contain  only  NH4+-  and  Ac~- 
ions,  H2O,  and  a  very  small  quantity  of  undissociated  NH4Ac.  It  is, 
then,  the  energy  consumed  in  the  dissociation  of  the  molecules,  as  indi- 
cated in  the  reactions  above,  which  has  lessened  the  total  heat  evolved 
on  neutralization  as  compared  with  the  neutralization  of  two  strong 
electrolytes,  which  are  almost  completely  dissociated  before  they  are 
brought  into  contact. 

The  neutral  solution  of  the  salt  ammonium  acetate,  since  it  contains 
a  far  greater  number  of  ions  than  either  the  base  or  acid  from  which  it 
is  formed,  will  also  be  a  far  better  conductor  of  electricity. 

THE  ACTION  OF  A  STRONG  ACID  UPON  THE  SALT  OF  A  WEAK  ACID 

41.  It  has  already  been  shown,  on  page  27,  that  when  an  insoluble 
substance  can  result  from  a  combination  of  ions  within  a  solution,  the 
reaction  continues  until  one  or  both  of  the  reacting  substances  is 
exhausted ;  and  it  has  also  been  demonstrated  in  the  preceding  para- 
graphs that  the  same  conditions  hold  when  H+-  and  OH  "-ions  unite  to 
form  water.  In  both  instances  the  product  is  an  undissociated  body. 

On  the  other  hand,  it  has  been  shown  (page  26)  that,  if  the  possible 
products  from  the  interaction  of  the  ions  are  all  substances  which  are 
largely  ionized,  no  evidence  of  chemical  change  is  manifest.  Between 
these  two  extremes  there  are  many  cases  in  which  the  product  is  ion- 
ized, but  only  to  a  moderate  extent,  and  in  such  instances  a  reaction 
will  occur,  but  it  will  only  partially  complete  itself.  An  example  is 
found  in  the  action  of  a  strong  acid,  as  hydrochloric  acid,  upon  the  salt 
of  a  weak  acid,  as  sodium  acetate.  If  these  substances  are  brought 


38  Electrolytic  Dissociation    Theory 

together  the  odor  of  undissociated  acetic  acid  (that  of  vinegar)  is  at 
once  noticeable.  This  is  occasioned  by  the  union  of  the  H+-ions  from 
the  ionized  hydrochloric  acid  with  the  Ac~-ions  of  ionized  sodium  acetate 
to  form  the  acetic  acid,  which  (see  Appendix)  is  relatively  little  disso- 
ciated. The  reaction  is  essentially  H+,  Cl"  -f-  Na+,  Ac" >  (HAc)  -f- 

Na"1",  Cl",  and  the  major  portion  of  the  weak  acetic  acid  will  be  dis- 
placed from  its  salt.  Since,  however,  a  portion  of  the  acetic  acid  is 
dissociated,  whereby  H"1"-  and  Ac~-ions  would  result,  it  is  evident  that 
the  reaction  cannot  continue  quite  to  completion. 

Thus,  any  strong  acid  will  displace  from  a  neutral  salt  of  a  weak 
acid  the  major  part  of  that  weak  acid  in  its  undissociated  form. 

42.  The  Effect  of  the  Formation  of  Volatile  Products.  — A  third 
case  remains  to  be  discussed,  namely,  that  in  which  one  of  the  products  of 
the  interaction  of  the  ions  is  relatively  volatile. 

Even  a  weaker  acid  may  displace  a  slight  amount  of  a  stronger  acid 
from  its  neutral  salt,  as,  for  instance,  when  sulphuric  acid  acts  upon  sodium 
chloride  a  certain  amount  of  undissociated  hydrogen  chloride  is  formed. 

If,  as  is  true  in  this  case,  the  slight  amount  of  the  new  substance  formed 
is  volatile,  its  escape,  by  reducing  its  concentration,  renders  the  formation  of 
fresh  quantities  possible,  and  under  the  proper  conditions  the  complete  dis- 
placement of  the  stronger  acid  may  be  effected  by  means  of  the  weaker  one. 
Use  is  made  of  this  fact  in  the  commercial  manufacture  of  hydrochloric  acid, 
in  which  common  salt  and  concentrated  sulphuric  acid  when  heated  together 
react  to  form  sodium  sulphate  and  gaseous  hydrogen  chloride,  the  latter 
of  which,  when  conducted  away  and  dissolved  in  water,  forms  the  common 
hydrochloric  acid  of  commerce.  This  is  explained  in  detail  as  follows :  At 
a  high  temperature  sodium  chloride  dissolves  to  some  extent  in  concentrated 
sulphuric  acid,  forming  Na+-  and  Cl~-ions;  concentrated  sulphuric  acid  is 
also  somewhat  dissociated,  with  the  formation  of  a  small  quantity  of  H+-ions. 
According  to  the  principle  of  Mass  Action  a  certain  proportion  of  the 
H+-ions  and  Cl~-ions  must,  under  these  conditions,  unite  to  form  undisso- 
ciated HC1 ;  but,  since  the  amount  of  the  latter  which  does  form  is  in  excess 
of  its  solubility  in  the  strong  sulphuric  acid,  it  escapes  as  a  gas  nearly  as 
fast  as  it  is  produced.  Since  the  liquid  will  not  hold  in  solution  sufficient 
undissociated  hydrogen  chloride  to  maintain  an  equilibrium  with  its  ioniza- 
tion  products,  the  latter  must  continuously  combine,  and  since  their  con- 
centration is  constantly  maintained  through  the  progressive  dissociation  of 
sulphuric  acid  and  sodium  chloride,  the  process  is  only  ended  when  either 
one  or  both  of  these  materials  is  used  up.  The  apparent  reaction  is :  — 

2NaCl   +  H.2S04  =  Na2S04  +  2HC1. 

It  is  the  volatility  of  hydrogen  chloride  which  makes  its  formation  possible ; 
for,  if  it  could  not  escape  from  the  reacting  mixture,  the  presence  of  a  very 


Solubility  of  Carbonates  39 

small  amount  of  it  would  prevent  its  further  formation,  since  it  is  a  stronger 
(/.  e.,  more  largely  ionized)  acid  than  sulphuric  acid. 

43.  In  general,  then,  a  reversible  reaction  can  continue  to  a  complete 
disappearance  of  one  or  both  of  the  reacting  substances  upon  one  side  of 
the  equation,  if  one  of  the  products  formed  upon  the  other  side  is  removed 
from  the  sphere  of  action  as  fast  as  it  is  produced.     This  occurs  if  this 
product  is  insoluble,  and  either  falls  as  a  precipitate  or  escapes  as  a  gas. 
In  addition,  as  has  already  been  explained  under  neutralization  of  acids  and 
bases,  reactions  may  continue  to  the  complete  disappearance  of  the  reacting 
bodies,  if  one  or  more  of  the  products  are  practically  undissociated  (as  is  the 
case  with  water),  even  when  these  products  do  not  escape  from  the  solution. 
This  is  one  of  the  most  important  consequences  of  the  Mass  Action  Law. 

44.  The    Solubility  of  Carbonates   in   Acids.  —  Most  carbonates, 
except  those  of  the  alkali  metals,  are  so  little  soluble  in  water  that  they  are 
generally  spoken  of  as  insoluble  compounds.     Thus,  if  solutions  of  sodium 
carbonate  and  calcium   chloride   are  poured  together,   a   precipitate   forms, 
according  to  the  reaction:  — 

Ca++,  Cr,  Cr,  +  Na+,  Na+,  CO8~~  =  2Na+,  2Cr  +  CaCO3. 

Calcium  carbonate  is,  however,  soluble  in  acids  with  an  evolution  of  carbon 
dioxide  gas,  as  was  noted  under  the  discussion  of  acids  as  one  of  the  char- 
acteristic effects  of  hydrogen  ions.  It  is  necessary,  in  order  to  explain  why 
the  presence  of  hydrogen  ions  should  cause  an  otherwise  insoluble  substance 
to  dissolve,  to  assume  that  calcium  carbonate  is  not  absolutely  insoluble  in 
water,  but  that  water  in  contact  with  it  dissolves  a  very  small  amount,  even 
although  this  amount  may  be  so  exceedingly  small  that  a  solution  saturated 
with  it  will  leave  no  perceptible  residue  after  evaporation.  This  very  small 
amount  which  is  dissolved  consists  of  Ca++-  and  CO3~~-ions.  Carbonic  acid, 
which  is  a  very  weak  acid,  dissociates  primarily  into  hydrogen  ions  a*nd 
complex  negative  ions,  H2CO3  ^  H+  -f-  HCO3~,  the  latter  of  which  dis- 
sociates further,  but  even  to  still  less  an  extent,  into  more  hydrogen  ions  and 
carbonate  ions,  HCO3~  ^  H+  -|-  CO3~.  Although  the  solubility  product  for 
Ca++-  and  CO3~~-ions  is  so  extremely  small,  nevertheless  that  for  Ca++-ions 
and  the  more  complex  HCO3~-ions  is  large ;  that  is  to  say,  calcium  acid 
carbonate,  Ca(HCO3)2,  is  a  very  soluble  salt,  as  well  as  one  which  is  to  a 
high  degree  ionized. 

When  powdered  calcium  carbonate  is  suspended  in  water,  it  dissolves 
until  the  solubility  product  of  Ca++-ions  and  CO3~~-ions  is  reached.  If  any 
acid,  such  as  hydrochloric  acid,  is  then  added,  the  CO3-ions  cannot  longer 
remain  in  even  their  former  small  concentration  in  presence  of  the  H+-ions 
of  the  acid ;  but  the  major  part  of  them  combine  to  form  HCO3~-ions,  and 
these  again  combine  with  more  H+-ions  to  form  the  H2CO3-molecules.  But 
by  this  action  the  product  of  the  concentrations  of  Ca"*"1"-  and  CO8~~-ions 
is  diminished  below  the  solubility  product,  so  that  more  calcium  carbonate  is 
able  to  dissolve.  This  process  is  a  continuous  one  so  long  as  there  remain 
H+-ions  in  the  solution  and  undissolved  calcium  carbonate ;  for,  long  before 
enough  carbonic  acid  can  have  formed  in  solution  to  cause  the  process  to 


40  Electrolytic  Dissociation   Theory 

reverse,  it  has  itself  decomposed  into  water  and  carbon  dioxide,  H2CO3 > 

H2O  +  CO2,  the  latter  escaping  as  a  gas. 

The  completeness  with  which  this  reaction  takes  place  is  due  to  two 
causes :  first,  to  the  slight  tendency  of  the  HCO3~-ions,  as  well  as  of  H2CO3- 
molecules,  to  dissociate;  and  second,  to  the  insolubility  of  carbon  dioxide  in 
water. 

45.  The  Precipitation  of  the  Metallic  Sulphides. — When  a  stream 
of  hydrogen  sulphide  gas  is  passed  into  an  acid  solution  of  copper  chloride, 
a  black  precipitate  of  copper  sulphide  is  formed.  Hydrogen  sulphide  is 
somewhat  soluble  in  water,  forming  in  it  a  very  weak  acid,1  H2S  ^  H+  -I'- 
ll+  -)-  S~  ~.  In  a  solution  containing  also  hydrochloric  acid,  the  concentra- 
tion of  S~~-ions  will  be  very  much  diminished  below  that  in  a  solution 
containing  no  strong  acid ;  for,  in  the  former  case,  the  concentration  of  the 
H+-ions  is  many  times  increased,  and  the  dissociation  of  the  hydrogen  sul- 
phide, and  consequently  the  concentration  of  the  S~~-ions,  must  be  corre- 
spondingly greatly  decreased  in  accordance  with  the  Law  of  Mass  Action. 
Therefore  the  solubility  product  of  copper  sulphide  must  be  exceedingly 
small,  if  even  in  a  strongly  acid  solution  of  hydrogen  sulphide  there  are 
sufficient  S~~-ions  to  form  a  precipitate  with  Cu++-ions. 

Zinc  behaves  differently  from  copper  in  that  it  is  only  partially  thrown 
down  in  the  form  of  sulphide  when  hydrogen  sulphide  is  passed  into  a 
neutral  solution  of  zinc  chloride,  or  one  faintly  acid  with  hydrochloric 
acid,  while  precipitation  is  entirely  prevented  by  any  considerable  quantity 
of  hydrochloric  acid.  An  explanation  of  this  is  found  in  the  fact  that  the 
solubility  product  of  zinc  sulphide,  although  small,  is  still  very  many  times 
greater  than  that  of  copper  sulphide.  As  long  as  the  solution  of  the  zinc 
salt  is  neutral  or  only  feebly  acid,  the  hydrogen  sulphide  is  sufficiently  dis- 
sociated to  yield  S~  "-ions  in  such  quantity  that  the  solubility  product  of 
the  zinc  sulphide  is  exceeded,  and  precipitation  ensues.  But  as  soon  as  the 
concentration  of  the  hydrochloric  acid  is  considerable,  the  total  concentration 
of  the  H+-ions  is  so  far  increased  that  that  of  the  S~~-ions  is  decreased  to 
a  point  below  that  at  which,  multiplied  by  the  concentration  of  the  Zn++-ions, 
it  gives  a  quantity  equal  to  the  solubility  product  of  zinc  sulphide,  therefore 
precipitation  can  no  longer  occur. 

That  the  solubility  product  of  ferrous  sulphide,  again,  is  greater  than  that 
of  zinc  sulphide  or  copper  sulphide  is  indicated  by  its  failure  to  precipitate 
when  hydrogen  sulphide  is  passed  into  even  a  neutral  solution  of  ferrous 
chloride.  If,  on  the  other  hand,  a  solution  of  ammonium  sulphide  is  added 
to  such  a  solution  the  iron  is  at  once  completely  precipitated  as  ferrous 
sulphide,  because,  ammonium  sulphide  being  a  salt,  it  is  largely  ionized  in 
solution,  and,  since  it  is  very  soluble  in  water,  it  yields  so  many  S"~-ions 
that,  even  with  a  small  amount  of  Fe++-ions,  the  solubility  product  of  the 
ferrous  sulphide  is  readily  exceeded.  For  the  same  reason  zinc  would  also 

1The  intermediate  ion,  HS~,  which  dissociates  further  into  H+-  and  S — -ions,  is  also 
present  here,  but  in  this  explanation  simplicity  is  gained,  and  no  serious  error  is  made,  by 
leaving  it  out  of  consideration. 


'HE  ACTION  OF  A  STRONG  BASE  UPON  THE  SALT  OF  A  WEAK  BASE 


f 
Neutral  Salts  and   Weak  Acids  41 

be  completely  precipitated  by  this  reagent  from  a  solution  of  a  zinc  salt. 
The  ionic  changes  in  the  case  of  the  iron  salt  are  as  follows :  — 

NH4+,  NH4+,  S~"  -f  Fe++,  Cl",  Cl"  =  FeS  +  2NH4+,  2Cr. 

It  should  be  remembered  that  the  solubility  product  of  ferrous  sulphide, 
although  larger  than  that  of  the  zinc  sulphide,  is  still  small,  otherwise  no 
precipitation  could  occur. 

" 

46.  If  sodium  hydroxide  is  added  to  a  solution  of  ammonium 
chloride  the  odor  of  ammonia  becomes  apparent.  The  OH  "-ions  of 
the  strong  base,  sodium  hydroxide,  unite  with  NH4+-ions  of  the  salt  to 
form  the  weakly  dissociated  base,  ammonium  hydroxide,  Na+,  OH~  -\- 
NH4+,  Cl-  =  (NH4OH)  +  Na+  Cl~  Ammonium  hydroxide  disso- 
ciates also  non-electrolytically  to  some  extent  into  water  and  ammonia, 
NH4OH  ^  H2O  +  NH3,  and  the  small  amount  of  the  latter  which 
escapes  as  a  gas  indicates  by  its  odor  the  existence  of  ammonium 
hydroxide  in  the  solution. 

The  action  of  a  strong  base  upon  a  salt  of  a  weak  base  results 
always,  as  in  the  above  instance,  in  the  formation  of  a  greater  or  less 
amount  of  the  undissociated  weak  base.  It  should  be  noted  that  this 
is  comparable  with  the  action  of  a  strong  acid  upon  the  salt  of  a  weak 
acid. 


EFFECT  UPON  THE  PROPERTIES  OF  WEAK  ACIDS  OR  BASES  OF  NEUTRAL 
SALTS  WITH  A  COMMON  ION 

47.  The  acid  or  basic  character  of  weak  acids  or  bases  is  very  greatly 
reduced  by  the  presence  of  any  highly  ionized  neutral  salt  with  a  common 
anion  or  cation.  This  will  be  clear  from  the  following  illustration  :  — 

In  a  solution  of  acetic  acid  a  comparatively  small  number  of  HAc-mole- 
cules  dissociate  into  an  equal  number  of  H+-  and  Ac~-ions,  until  a  state  of 
equilibrium  is  reached,  when 

Contn  H+  X  ConSn  Ac~ 

-— — — — =  Constant. 

Concn  HAc 

In  a  molal  solution  of  pure  acetic  acid  the  H+-ions  and  Ac~-ions  have  been 
proved  to  be  present  in  a  concentration  0.0042  molal,  *.  e., 

.0042  X  .0042 


•9958 


=  .000018. 


42  Electrolytic  Dissociation   Theory 

But  in  a  solution  which  is  molal  with  regard  to  sodium  acetate  as  well  as 
acetic  acid  (since  sodium  acetate  at  that  concentration  is  known  to  be  50  per 
cent,  dissociated),  the  concentration  of  Ac~-ions  will  be  0.50  molal.  The 
concentration  of  undissociated  acetic  acid  molecules  is  so  nearly  that  of  the 
total  amount  present  (i.o  molal)  that  it  may  practically  be  placed  at  that 
figure.  Calling  x  the  concentration  of  H+-ions,  the  formula  becomes  :  — 

x  X  0.150 

—  =  0.000018,     or     x  =  0.00004. 

I.O 

Therefore  the  concentration  of  the  H+-ions  has  been  reduced  to  T£Q  of  its 
previous  value  by  the  presence  of  an  equivalent  quantity  of  neutral  sodium 
acetate,  or,  to  put  it  in  another  form,  the  strength  or  efficiency  of  the  acetic 
acid  has  been  lessened,  since,  as  has  already  been  stated  (page  33),  the 
strength  of  an  acid  depends  upon  the  concentration  of  the  H+-ions  which 
it  yields. 

The  basic  character  of  ammonium  hydroxide  is  so  diminished  by  the 
presence  of  ammonium  chloride  that,  as  will  be  mentioned  in  Chapter  V, 
ammonia  fails,  in  the  presence  of  the  latter  reagent,  to  give  the  usual 
precipitates  of  the  hydroxides  of  cobalt,  magnesium,  etc. 

HYDROLYSIS 

48.  When  perfectly  pure  potassium  cyanide  is  dissolved  in  water 
the  solution  acquires  an  alkaline  reaction  ;  that  is,  it  exhibits  the  char- 
acteristic reactions  for  OH~-ions,  and  turns  red  litmus  blue.  When 
pure  ferric  chloride  is  dissolved  in  water  the  solution  acquires  an  acid 
reaction  ;  that  is,  it  shows  the  reactions  of  H+-ions,  and  turns  blue 
litmus  red. 

In  order  to  explain  these  phenomena  it  is  necessary  to  take  into 
account  certain  properties  of  the  pure  solvent  water  which  have  been 
disregarded  in  the  preceding  discussions,  namely,  the  presence  in  the 
water  of  a  very  small  quantity  of  H+-  and  OH~-ions.  It  is  quite 
true  that  the  extent  of  this  dissociation  is  so  slight  that  it  may 
ordinarily  be  neglected  —  that  it  does  not  appreciably  affect,  for  ex- 
ample, the  completeness  of  the  neutralization  reactions  of  strong  acids 
and  bases.  On  the  other  hand,  it  does  attain  some  significance  when 
it  becomes  necessary  to  consider  bodies  which,  while  exhibiting  greater 
dissociation  than  the  water,  are  not  far  removed  from  it  in  this  particu- 
lar.1 In  such  cases  the  phenomena  known  as  hydrolysis  makes  itself 


1  In  i  liter  of  pure  water  there  is  present  but  To^oV^o  o  of  a  mo1  of  H"*"'  or  OH~-ions. 
The  extent  to  which  some  of  the  most  important  electrolytes  are  dissociated  is  given  in 
the  Appendix. 


Hydrolysis  43 

evident ;  that  is,  the  dissociated  water  takes  a  part  in  the  chemical 
action.  This  will  be  best  understood  by  considering  in  detail  the 
hydrolysis  of  the  two  salts  mentioned  above. 

These  two  substances,  in  conformity  with  the  general  rule  regard- 
ing neutral  salts,  are  highly  ionized  :  — 

KCN  -   ->K+-f-Cn-; 

FeC]3  -  -*  Fe+  +  +  +  CT  +  CT  +  Cr. 

Hydrocyanic  acid,  however,  which  is  the  acid  from  which  potassium 
cyanide  is  formed,  is  a  substance  which  is .  dissociated  but  little  more 
than  water,  and  although  the  amount  of  H+-  and  OH  "-ions  into  which 
pure  water  is  dissociated  is  so  slight,  yet  the  degree  of  dissociation  of, 
hydrocyanic  acid  is  also  so  small  that  not  even  the  H+-ions  normally 
present  in  water  can  remain  as  such  in  the  presence  of  the  CN~-ions  of 
potassium  cyanide.  A  formation  of  undissociated  hydrogen  cyanide 
results,  H+  +  CN~  — >  HCN.  Since  the  equilibrium  between  water 
and  its  ions  is  thereby  disturbed,  more  water  must  dissociate  tD  fur- 
nish a  fresh  supply  of  H+-ions,  H2O  -  -f>  H+  +  OH~.  The  H+-ions 
thus  formed  being  once  more  in  excess  of  the  number  which  can 
remain  in  equilibrium  with  the  CN~-ioris  and  the  HCN-molecules 
already  formed,  the  process  continues  until  a  point  of  equilibrium  is 
reached  when  the  amount  of  undissociated  hydrogen  cyanide  no  longer 
increases. 

Since  the  progressive  dissociation  of  water  must  have  yielded  OH~- 
ions  equal  in  number  to  the  H+-ions,  and  since  those  ions  have  but  a 
very  slight  tendency  to  combine  with  K+-ions  to  form  undissociated 
KOH,  they  must  have  accumulated  in  the  solution,  thereby  giving  it 
its  basic  character. 

The  hydrolysis  of  potassium  cyanide  thus  consists  of  a  partial  reac- 
tion of  this  salt  with  water,  with  a  formation  of  equivalent  amounts  of 
free  hydrocyanic  acid,  undissociated  (and  therefore  inactive),  and  of  free 
potassium  hydroxide  (dissociated  and  active),  the  latter  of  which  gives 
to  the  solution  its  basic  properties.  That  hydrolysis  is  exactly  the 
reverse  of  neutralization  is  shown  by  this  equation  :  — 

HCN  +  K+,  OH'  ^  HOH  +  K4,  CN~. 


44  Electrolytic  Dissociation   Theory 

\ 

49.  The  hydrolysis  of  ferric  chloride,  on  the  other  hand,  can  only 
result  in  the  formation  of  hydrochloric » acid  and  ferric  hydroxide,  the 
latter  a  body  which  is  about  as  little  dissociated  as  hydrogen  cyanide. 
Hence  in  a  solution  of  ferric  chloride  the  OH  "-ions,  which  are  normally 
present  in  pure  water,  cannot  exist  in  the  presence  of  the  Fe+  +  +-ions, 
and  a  formation  of  undissociated  ferric  hydroxide  is  the  result,  Fe+  +  +  -f- 

3OH~ >  Fe(OH)3.     This  removal  of  OH~-ions  permits  more  water 

to  dissociate,  and  this  process  continues  until  an  appreciable  amount 
of  undissociated  ferric  hydroxide  and  of  ionized  hydrochloric  acid  has 
been  formed,  the  H+-ions  of  the  latter  imparting  to  the   solution  its 
acid  reaction.     The  hydrolysis  of  ferric  chloride  is  thus   exactly  the 
reverse  of  the  neutralization  of  ferric  hydroxide  and  hydrochloric  acid  :  — 

Fe03H3  +  3H+,  3C1-  ^  Fe+  +  +,  3CT  +  3H2O. 

50.  It  is  to  be   noted  that  in  the   case   of  potassium   cyanide  a 
weak  acid  is  combined  with  a  strong  base,  while  in  ferric  chloride  the 
reverse  is  true.     Salts  formed  from  strong  acids  and  strong  bases  are 
not  hydrolyzed  and  give  a  perfectly  neutral  reaction  with  litmus.     The 
reason  for  this  will  be  evident  if  it  is  recalled  that  when  sodium  chloride, 
for  example,  is  dissolved  it  dissociates  into  Na+-  and  Cl~-ions,  neither 
of  which  have  more  than  a  very  slight  tendency  to  combine  with  OH~- 
or  H+-ions,  respectively,  and  that  therefore  the  amounts  of  undissociated 
NaOH  and  HC1  which  are  formed  are  entirely  too  small  to  disturb  the 
equilibrium  which  exists  between  water  and  its  ions. 

On  the  other  hand,  the  hydrolysis  of  a  salt  which  is  formed  from 
an  acid  and  a  base  both-  of  which  are  weak  may  be  so  complete  that 
it  will  not  be  possible  for  the  salt  to  exist  in  solution.  For  example, 
when  ammonium  sulphide  is  added  to  a  solution  of  aluminum  chloride, 
instead  of  the  formation  of  aluminum  sulphide  within  the  solution,  there 
will  appear  its  hydrolyzation  products,  aluminum  hydroxide  and  hydro- 
gen sulphide.  Since  both  of  these  will  be  formed  in  excess  of  their 
solubility,  the  former  will  fall  as  a  precipitate  and  the  latter  will  escape 
as  a  gas  :  — 

+,  6C1-  +  6NH4+,  38--  +  6H+,  6(OH)-  = 
6NH4+,  6C1-  +  2A1(OH)3  +  3(H2S). 


CHAPTER  III 

ELECTROLYTIC  SOLUTION  PRESSURE 

51.  It  is  well  known  that  where  strips  of  zinc,  iron  nails,  or  other 
fragments  of  iron  are  immersed  in  solutions  of  copper  salts,  the  zinc  or 
the  iron  at  once  becomes  coated  with  copper,  and  a  qualitative  exami- 
nation of  the  solution  shows  that  some  of  these  metals  have  simultane- 
ously dissolved.  This  process  of  interchange  will  continue  until  the 
zinc  or  iron  has  entirely  dissolved,  leaving  in  its  place  a  spongy  mass 
of  metallic  copper ;  or,  if  the  zinc  or  iron  is  in  excess,  the  process  will 
continue  until  all  of  the  copper  is  removed  from  the  solution.  Platinum 
or  silver,  on  the  other  hand,  if  immersed  in  a  solution  of  the  same 
copper  salt  causes  no  visible  change.  To  explain  these  differences  it 
is  necessary  to  consider  in  detail  the  process  which  has  gone  on  between 
the  zinc  and  the  copper  salt. 

It  is  found  by  quantitative  experiments  that  for  every  63.6  grams 
of  copper  deposited  65.4  grams  of  zinc  pass  into  solution.  These 
quantities  represent  the  atomic  weights  in  grams  of  the  two  metals. 
It  is,  therefore,  evident  that  for  every  atom  of  copper  deposited  from 
solution  an  atom  of  zinc  passes  into  solution ;  and,  with  this  as  a 
starting  point,  the  behavior  of  the  two  metals  is  easily  explained  by 
the  ionic  theory,  as  follows  :  The  metallic  zinc  is  made  up  of  electrically 
neutral  zinc  atoms ;  copper  sulphate  in  solution  exists  essentially  in  the 
form  of  Cu++-  and  SO4~~-ions ;  zinc  sulphate  exists  in  solution  as  Zn++- 

and  SO4~~-ions.  The  real  change,  then,  is,  Zn  -f-  Cu++,  SO4~~ > 

Cu-|-Zn++,  SO4~~;  that  is,  an  electrically  neutral  zinc  atom  has 
become  a  Zn++-ion  by  assuming  positive  charges,  while  a  Cu++-ion 
has  become  an  electrically  neutral  copper  atom  by  giving  up  exactly 
equivalent  positive  charges.  This  expressed  as  an  equation  is :  — 

Zn-f  Cu++  =  Cu  +  Zn++. 

In  still  greater  detail  this  process  may  be  described  as  follows :  The  fact 
that  the  strip  of  zinc  is  electrically  neutral  implies  that  whatever  charges  of 

45 


46  Electrolytic  Dissociation   Theory 

positive  and  negative  electricity  reside  upon  it  must  be  equal.  As  the 
Zn++-ions  separate  from  it,  carrying  their  positive  charges,  the  strip  retains 
residual  negative  charges  which  attract  the  positively  charged  Cu++-ions  of 
the  solution,  which  thus  have  their  charges  neutralized  by  the  negative 
charge  on  the  strip  of  zinc,  and  deposit  as  copper  atoms.  The  solution 
pressure  of  copper  is  so  small  that  these  atoms  do  not,  like  those  of  zinc, 
tend  to  return  to  the  solution  as  ions. 

The  metals  cannot  dissolve  in  the  form  of  neutral  atoms,  but  must 
enter  the  solution  as  electrically  charged  ions ;  arid  experiment  has 
shown  that  every  metal  has  a  definite  tendency  to  pass  from  the 
atomic  to'  the  ionic  Condition,  just  as  the  zinc  does  in  the  above  illus- 
tration. This  impelling  force  is  known  as  the  electrolytic  solution 
pressure  for  the  metal  in  question,  and  this  quantity  is  found  to  be 
capable  of  estimation  and  to  be  a  characteristic  property  of  the  metal. 

It  is  evident  that  the  metal  zinc  in  the  illustration  above  cannot,  in  spite 
of  the  great  solution  tension  which  it  possesses,  simply  pass  into  solution 
as  ions  with  no  other  accompanying  change ;  for  if  this  did  occur,  even  to 
a  very  slight  extent,  the  solution  would  then  be  charged  with  a  great  sur- 
plus of  positive  electricity,  since  the  ions  bear  enormous  charges.  The 
'zinc  ions,  with  their  positive  charges,  would  now,  not  being  balanced  by 
negatively  charged  ions,  repel  each  other  with  a  force  far  greater  than  the 
.solution  tension  of  the  .metal,  so  that  the  zinc  ions  would  be  forced  out 
of  solution  again  into  the  original  neutral  metallic  condition. 

But  if  the  strip  of  zinc  is  placed  in  a  solution  containing  positive  ions 
of  some  element  with  a  lesser  solution  tension  (see  table  on  page  47),  as,  for 
example,  the  solution  of  copper  sulphate,  there  will  then  be  an  outlet  by  which 
the  positive  electricity  carried  into  solution  by  the  zinc  ions  can  escape  — 
or,  more  exactly,  by  which  an  equivalent  amount  of  positive  electricity  may 
escape — •  namely,  by  the  giving  up  of  the  positive  charges  of  the  copper 
kms  at  the  surface  of  the  zinc,  whereby  they  themselves  assume  the  form 
of  neutral  copper  atoms,  and  their  positive  electricity  neutralizes  the  nega- 
tive electricity  induced  on  the  metal  strip  as  the  result  of  the  formation  of 
positive  ions. 

52.  The  electrolytic  solution  pressure  may  be  regarded  as  a  phe- 
nomenon of  the  same  general .  nature  as  osmotic  pressure ;  in  fact,  the 
force  of  this  tension  may  be  expressed  in  terms  of  the  osmotic  pressure 
of  the  ions  of  the  respective  metals,  which  will  just  suffice  to  counter- 
balance it.  In  the  table  below  the  metals  are  arranged  in  the  order 
of  the  magnitude  of  their  respective  solution  pressures,  the  values  of 
which  are  expressed  in  atmospheres-  so  far  as  they  are  accurately  known. 


Solution   Pressure 


47 


ELECTROLYTIC  SOLUTION  PRESSURE  OF  THE  METALS  * 


Na+ 

Lead    Pb+  + 

1.1  X  10-8 
9.0  X  10-4 

4.8  X  10"* 
1.1  X  10~16 
2.3  X  10~17 
1.5  X  10~36 

Calcium 

Hydrogen    .    .    ,  H  + 
Bismuth  ....  Bi-*--1--1- 
Antimony    '. 

Magnesium 
Aluminum 
Manganese 
Zinc     .    .    . 
Cadmium    . 

.    .    .  Mg+  +           1.1  X  1043 
.    .    .  A1+  +  +               .    .    . 

.    .    .  Zn+  +            9.9  X  1018 
.    .    .  Cd++           2.7  X  106 
Fe++            1  2  X  104 

Arsenic 

Copper    .    .    .    .  Cu+  + 
Mercury  ....  Hg+ 
Silver  -^R"*" 

Cobalt 

Co+  +           1  9 

Palladium    .    .    .  Pd+  + 
Platinum          .        ... 

Nickel 

Ni+  +            1  3 

Tin 

Sn++                     .    . 

Gold     

The  electro-positive  character  of  the  metals  decreases  in  the  same 
order  as  the  solution  pressure,  and,  in  general,  any  metal  if  immersed 
in  the  solution  of  a  salt  of  a  metal  with  a  less  solution  pressure  (i.  e., 
one  which  is  less  electro-positive)  tends  to  cause  the  deposition  of  that 
metal  from  solution  as  the  zinc  or  iron  precipitates  copper. 

Certain  apparent  exceptions  to  this  principle  are  referred  to  below. 

53.  It  will  be  noted  that  hydrogen  has  been  included  in  the  table 
above.  This  element  resembles  the  metals  in  possessing  a  definite  solu- 
tion pressure  ;  and  since  hydrogen  ions  are,  as  we  have  seen,  always 
present  in  minute  quantities  in  aqueous  solutions  and  in  considerable 
quantities  in  acid  solutions,  the  relation  of  this  pressure  to  those  of  the 
various  metals  becomes  of  importance. 

All  metals  with  a  greater  solution  tension  than  hydrogen  will  dis- 
solve in  acids  with  the  liberation  of  hydrogen  gas,  provided  the  salts 

1  The  position  of  those  metals  for  which  no  value  is  given  is  only  approximate.  The 
values  which  are  given  represent  the  tendency  to  form  the  particular  ion  of  the  metal 
which  is  indicated.  Those  metals  which  are  not  followed  by  the  symbol  of  an  ion  do  not 
possess,  to  more  than  an  infinitesimal  extent,  the  power  to  yield  simple  positive  ions  in 
solution. 

It  will  be  noticed  that  copper  is  placed  above  mercury  and  silver  in  the  series,  although 
it  has  a  smaller  solution  pressure  than  either  of  those  metals.  This  is  because  the  effect 
of  the  double  positive  charge  on  the  Cu++-ion  more  than  counterbalances  the  greater 
solution  pressure  of  the  mercury  or  the  silver,  so  that  copper  is  able  to  replace  these 
metals  from  their  solutions. 


48  Electrolytic  Dissociation   Theory 

formed  are  not  themselves  insoluble.  The  action  of  hydrochloric  acid 
upon  zinc  is  a  familiar  instance,  and  demands  consideration  in  detail. 
When  a  piece  of  zinc  is  immersed  in  the  acid  its  solution  pressure  at 
once  causes  it  to  send  off  zinc  ions  into  the  solution,  and  as  these  ions 
with  their  positive  charges  leave  the  mass  of  zinc  the  latter  becomes 
negatively  charged,  as  described  at  the  top  of  page  46.  The  H+- 
ions  in  the  acid  solution,  like  the  Cu++-ions  in  a  solution  of  a  copper 
salt,  now  discharge  themselves  upon  the  zinc,  become  hydrogen  atoms, 
and  these  combine  to  form  molecules  of  hydrogen  gas,  which  escape 
from  the  solution,  Zn  -f  2H+  =  (H2)  -+-  Zn++.  If,  on  the  other  hand, 
the  solution  pressure  of  the  metal  is  less  than  that  of  hydrogen,  as  is 
the  case  with  copper,  the  hydrogen  ions  are  not  discharged,  and  no 
apparent  action  ensues. 

54.  It  happens  in  some  cases  in  actual  practice  that  the  phenomena  as 
above  described  do  not  appear  to  take  place  according  to  the  theory ;  that 
is,  that  metals  with  lower  solution  pressure  are  not  always  thrown  out  of 
solution  by  those  of  higher  solution  tension.      For  example,  if  a  piece  of 
perfectly  pure  zinc  is  placed  in  a  solution  of  an  acid  there  is  no  evolution 
of  hydrogen  gas.     This  may  be  explained  by  assuming  that  a  certain  num- 
ber of  zinc  ions  do  go  into  solution,  thereby  forcing  an  equal  number  of 
hydrogen  ions  into  the  condition  of  uncharged  atoms,  but  that  these  atoms, 
instead  of  forming  molecular  hydrogen,  are  absorbed  uniformly  by  the  uni- 
form surface  of  the  pure  zinc,  and  thus  form  a  sort  of  protective  layer  so 
that  further  action  is  stopped.     If  now  a  piece  of  platinum  wire  is  brought 
into  contact  with  the  piece  of  zinc  beneath  the  liquid,  action  immediately 
commences.     Zinc  goes  into  solution  as  Zn+  +-ions  and  hydrogen  gas  escapes, 
but  only  from  the  surface  of  the  platinum  wire  and  not  from  the  zinc.     The 
platinum  has  broken  the  protective  coating  about  the  metal,  and,  being  in 
electrical  contact  with  it,  a  part  of  the  negative  charge  induced  by  the  loss 
of  positive  ions  flows  into  the  wire  and  attracts  the  positively  charged  H+- 
ions  which  are  here  free  to  discharge  themselves,  thus  leaving  the  solution 
in  a  condition  in  which  it  can  receive  more  positive  ions  from  the  zinc. 

In  ordinary  commercial  zinc  it  is  to  be  noted  that  the  surface  is  covered 
with  minute  particles  of  various  impurities  which  serve  the  same  end  as  the 
platinum  wire  above ;  that  is,  furnish  points  at  which  the  hydrogen  gas  can 
form  and  escape.  Commercial  zinc,  therefore,  as  is  well  known,  dissolves 
with  ease  in  acids. 

55.  Again,  it  is  to  be  noted  that  the  readiness  with  which  ions  of  a 
metal  pass  into  a  solution  is  influenced  by  the  number  of  ions  of  that  metal 
already  in  the  solution.     Thus  there  is  more  tendency  for  ions  of  cobalt  to 
leave  a  rod  of  cobalt  immersed  in  a  solution  containing  no  cobalt  ions  than 
if  the  solution  already  contained  a  cobalt  salt.     Conversely,  in  an  acid  solu- 
tion the  tendency  of  hydrogen  ions  to  pass  into  the  electrically  neutral  state 
is  greater  the  higher  the  concentration  of  these  ions.     It  is  for  this  reason 


lonization  Pressure  49 

that  metals  react  more  vigorously  the  stronger  the  acids  with  which  they  are 
in  contact,  because  the  reaction  is  a  result  of  the  combined  tendencies  of  the 
metal  to  pass  into  the  ionic  form  and  of  the  hydrogen  ions  to  pass  into 
the  neutral  condition.  Since  pure  water  contains  so  exceedingly  few  hydro- 
gen ions,  it  is  apparent  why  many  metals  which  possess  much  higher  solution 
tensions  than  hydrogen,  and  accordingly  attack  acids  readily,  have  no  effect 
upon  water. 

THE   ELECTROLYTIC   IONIZATION    PRESSURE   OF  THE   NEGATIVE   ELE- 
MENTS, OR  THE  TENDENCY  OF  THESE  ELEMENTS  TO  PASS 
INTO  THE  IONIC  CONDITION 

56.  When  chlorine  gas  is  bubbled  into  a  solution  of  potassium 
iodide,  free  iodine  is  liberated  which  colors  the  solution  brown,  and  the 
chlorine  is  found  to  have  taken  the  place  of  iodine  in  the  compound, 
forming  potassium  chloride  :  (C12)  -\-  2K+,  2\~  =  2K+,  2d~  +  (I2)  ;  or, 
more  simply,  (C12)  +  2l~  =  2C1~  +  (I2). 

This  case  differs  from  that  of  the  replacement  of  copper  by  zinc 
simply  in  the  fact  that  here  the  electrically  neutral  elements  are  them- 
selves soluble,  and  that  as  ions  they  assume  negative  instead  of  positive 
charges. 

Each  of  the  negative  elements  which  can  form  simple  ions  is 
characterized  by  a  certain  definite  tendency  to  do  so,  which  is  similar 
to  the  electrolytic  solution  tension  of  the  positive  elements,  and  they 
may  also  be  arranged  in  a  series  according  to  the  magnitude  of  their 
ionization  tensions.  Any  element  in  such  a  series  would  be  more 
electro-negative  than  any  element  standing  below  it,  and  would  dis- 
place the  latter  from  its  salts  when  dissolved  in  water. 

Although,  owing  to  the  experimental  difficulties  involved,  exact 
values  have  not  been  obtained,  yet  the  order  in  which  the  negative 
elements  stand  with  respect  to  their  tendency  to  form  ions  is  shown 
in  the  following  table  1 :  — 

Fluorine  .  .- '   •   .         .         .         .         .         .  F~ 

Chlorine  .  .         .         .         .         ...  Cl~ 

Bromine  .  .         .         .....         .  Br~ 

Oxygen  .  .         .... 

Iodine  .  .         .         .         .....  I~ 

Sulphur  .  .         . '       .      '    .         .         .         .  S~- 

1  Oxygen  would  naturally  first  form  the  ion  O  ;  this  ion  cannot,  however,  attain  any 
appreciable  concentration,  because  as  fast  as  it  is  formed  it  reacts  with  other  constituents 
of  the  solution  as  follows  :  O" -f  H+ =  OH~,  or  O"  +  H2O  =  2OH~. 


50  Electrolytic  Dissociation   Theory 

57.  It  has  already  been  stated,  on  page  18,  that  the  explanation  there 
given  in  connection  with  the  electrolysis  of  potassium  sulphate  did  not  take 
into  account  all  of  the  factors  involved.  It  has  since  been  shown  that  water 
is  itself  dissociated  to  a  slight  extent  into  H+-  and  OH~-ions,  and  it  may 
now  be  stated  that  an  exceedingly  small  fraction  of  the  latter  are  further 
dissociated  into  H+-  and  O~~-ions.  In  the  light  of  these  facts  the  full 
explanation  takes  the  following  form :  At  the  cathode  the  K+-ions  which 
bring  the  current  through  the  solution  to  that  electrode  might  give  up  their 
charges,1  but,  since  potassium  has  such  an  enormously  greater  solution  ten- 
sion than  hydrogen,  the  few  H+-ions,  in  the  layer  of  solution  immediately 
surrounding  the  cathode  will  give  up  their  charges  and  go  into  the  neutral 
condition  in  preference  to  the  K+-ions.  As  fast  as  the  H+-ions  disappear 
in  this  manner  more  water  molecules  will  dissociate  in  accordance  with  the 
Mass  Action  Law,  to  furnish  a  new  supply  of  H+-ions.  This  process  is  con- 
tinuous, and  the  OH~-ions,  which  are  also  formed  by  the  dissociation  of 
water,  remain  in  the  solution  and  electrically  balance  the  K+-ions  which  do 
not  discharge. 

At  the  anode  the  SO4~~-ions  which  bring  the  current  to  the  electrode 
might  discharge,  if  no  negative  ions  were  present  which  could  more  easily 
part  with  their  charges.  The  O~  "-ions  in  the  layer  surrounding  the  anode 
can,  however,  discharge  more  readily  than  the  SO4~~-ions,  because  their 
tendency  after  they  are  discharged  to  reassume  the  ionic  form  is  less ;  there- 
fore it  is  the  O~~-ions  which  do  give  up  their  negative  charges,  and  more 
water  molecules  dissociate  to  give  fresh  O~~-ions,  while  H+-ions  remain 
in  the  solution  to  electrically  balance  the  SO4~  ~-ions  which  do  not  discharge. 

1  Indeed,  this  does  occur  with  a  very  heavy  current  when  a  cathode  of  mercury  is 
used,  which  dissolves  the  potassium  set  free,  and  thus  lessens  its  tendency  to  react  with 
the  water.  -'  iX  <_ 


CHAPTER    IV 

K!K 

OXIDATION    AND    REDUCTION 

58.  When  iron  rusts  in  the  air,  forming  iron  oxide,  there  is  a  direct 
'addition  of  oxygen,  and  the  process  is  called  oxidation.     If  the  oxide 

is  heated  in  a  stream  of  hydrogen  and  its  oxygen  thereby  removed,  it  is 
said  to  be  reduced.  Instead  of  oxygen,  any  electro-negative  element,  as 
sulphur  or  chlorine,  can  be  made  to  unite  with  iron  to  form  iron  sul- 
phide or  iron  chloride,  and,  again,  by  suitable  means  iron  may  be  freed 
from  the  negative  element  in  these  compounds.  The  meaning  of  the 
terms  oxidation  and  reduction  has  been  extended  to  include  cases  such 
as  the  latter,  in  which  other  negative  elements  than  oxygen  are  added 
or  removed.- 

59.  When  iron  is  dissolved  in  hydrochloric   acid  it  passes  from 
the  electrically  neutral  metallic  condition  into  the  form  of  electrically 
charged  ions,  while  hydrogen  ions  simultaneously  release  their  charges, 
thereby  assuming  the  form  of  the  neutral  element.     If  the  resulting 
solution  is  evaporated,  solid  ferrous  chloride  is  obtained.     The  iron  is 
evidently  in  the  same  condition  as  if  it  had  been  made  to  combine 
directly  with  chlorine  gas ;  that  is,  it  is  in  the  oxidized  state.     It  may 

'be  asked,  Is,  then,  the  mere  fact  that  it  is  combined  with  a  negative 
"element  a  criterion  of  its  oxidation,  or  is  it  to  be  considered  as  oxidized 
even  when  in  solution  as  the  ion  Fe+  +  ?  The  important  change  which 
takes  place  during  the  transformation  of  iron  into  solid  ferrous  chloride, 
when  this  is  accomplished  by  first  dissolving  in  hydrochloric  acid  and 
•then  evaporating,  really  occurs  when  iron  atoms  acquire  positive  charges 
at  the  expense  of  hydrogen  ions ;  the  mere  gradual  evaporation  of  water, 
so  that  the  Fe++-  and  Cl~-ions  unite  more  and  more  as  FeCl2-molecules, 
does  not  essentially  alter  any  characteristic  of  ferrous  chloride.  There- 
fore the  oxidation  of  iron  when  it  is  dissolved  must  consist  in  the 
imparting  to  it  of  the  positive  charges  of  electricity,  and  the  substance 
from  which  it  receives  these  charges,  that  is,  hydrogen,  must  be 
correspondingly  reduced: 

51 


52  Electrolytic  Dissociation   Theory 

60.  But  iron  can  exist  in  different  states  of  oxidation.     If  chlorine 
gas  is  bubbled  into  the  solution  of  ferrous  chloride,  as  formed  above, 
until  no  more  is  absorbed,  and  this  solution  is  then  evaporated,  solid 
ferric  chloride,  FeCl3,  is   obtained.      The  iron  has  thus  been  further 
oxidized.     The  reaction   in   the  solution  is  as  follows :   2  Fe"1" +  -|-  4-Cl~ 
-|-  (C12)  =  2Fe++  '  -f-  6C1~.     The  ferrous  ion  has  acquired  one  more 
positive  charge,  by  which  it   is   changed  to  the  ferric  ion,  possessing 
properties  which  are  as  much  at  variance  with  those  of  ferrous  ions  as 
if  it  were  the  ion  of  an  entirely  different  element. 

But,  since  the  iron  has  been  oxidized,  there  must  have  been  some 
corresponding  reduction.  Chlorine  gas,  when  it  is  first  dissolved  in 
water,  does  not  dissociate,  but  forms  electrically  neutral  molecules. 
The  fact  that  the  molecules  are  electrically  neutral  implies  merely  that 
they  possess  no  excess  of  either  positive  or  negative  electricity.  They 
probably  do  possess  a  certain  quantity  of  each,  but  in  exactly  equal 
amounts.  At  the  moment  in  which  chlorine  takes  part  in  the  reaction 
noted  above  each  molecule  changes  into  two  negative  ions,  and  since 
two  negative  charges  have  thereby  appeared  the  two  positive  charges 
by  which  they  were  previously  held  inactive  must  have  also  appeared 
at  some  point.  They  have  attached  themselves  to  two  Fe++-ions,  con- 
verting them  into  Fe"1" +  +-ions,  which  are  now  capable  of  balancing  two 
additional  negative  Cl~-ions.  In  this  way  the  electrically  neutral  chlo- 
rine molecule  by  changing  into  chlorine  ions  has  lost  positive  charges, 
which  is  another  way  of  saying  that  it  has  undergone  reduction.  On 
the  other  hand,  the  Fe++-ions  have  acquired  positive  charges  at  the 
expense  of  the  chlorine  and  are  oxidized. 

61.  An  element,  then,  as  has  been  shown  for  iron,  can  be  oxidized 
(i)  by  causing  atoms  of  a  negative  element  to  combine  with  it ;  or  (2) 
by  imparting  positive  charges  of   electricity  to  its  atoms,  converting 
them  into  ions ;  or  (3)  by  imparting  additional  positive  charges  to  its 
already  existing  ions.     Ions,  however,  like  the  neutral  elements,  may 
be  oxidized  by  the  addition  of  atoms  of  negative  elements  as  well  as 
by  the  imparting  of  positive  charges,  as  illustrated  by  the  change  of 
SO3~~  to  SO4~~,  or  of  AsS3~~  "  to  AsS4        .     For  example,  when  a 
solution  of  sulphurous  acid  is  exposed  to  the  air,   it  slowly  absorbs 
oxygen,  and  sulphuric  acid  is  formed,   2H2SO3  +  O2  —  2H2SO4,  or 


Oxidation  and  Reduction  53 

SO3~"  +  O  =  SO4~~.  When  a  solution  of  sodium  sulpharsenite  is 
treated  with  sulphur  (dissolved  in  Na2S)  it  is  oxidized  to  sodium  sulph- 
arsenate,  Na3AsS3  +  S  =  Na3AsS4,  or  AsS3~ "  '  +  S  =  AsS4  . 

77/6'  oxidation  of  any  body  may,  then,  consist  in  the  addition  of 
atoms  of  a  negative  element  to  its  molecules,  atoms,  or  ions,  or  the  with- 
drawal of  the  atoms  of  a  positive  element ;  or  it  may  consist  in  the 
addition  of  positive  charges  of  electricity,  or  the  withdrawal  of  negative 
charges. 

Reduction  is  the  reverse  of  this,  namely,  the  addition  of  the  atoms  of 
positive  elements  or  of  negative  electrical  charges,  or  the  withdrawal 
of  the  atoms  of  negative  elements,  or  of  positive  electrical  charges. 

The  changes  in  the  state  of  oxidation  which  the  elements  may 
undergo  are  very  characteristic  of  them,  and,  like  the  formation  of 
precipitates  by  the  uniting  of  oppositely  charged  ions,  may  be  used 
extensively  in  the  identification  of  the  various  elements  in  qualitative 
analysis. 

In  Chapter  V,  under  the  reactions  of  the  ions,  will  be  found 
the  more  important  oxidation  changes  which  the  common  ions  may 
undergo,  but  in  the  next  few  paragraphs  some  typical  instances  will 
be  discussed  in  more  detail. 

62.  Tin  can  exist  in  its  compounds  in  two  different  states  of 
oxidation  which  correspond  to  the  oxides  SnO  and  SnO2,  in  which 
the  valence  of  tin  is  II  and  IV,  respectively.  Corresponding  to  these 
variations  in  valence,  tin  can  form  Sn++-  and  Sn+H  h+-ions,  which  exist, 
respectively,  in  solutions  of  stannous  and  stannic  chloride,  SnCl2  and 
SnCl4. 

Mercury  in  its  compounds  can  likewise  exist  in  two  different  states 
of  oxidation,  which  correspond  to  the  oxides  Hg2O  and  HgO,  in  which 
the  valence  is  I  and  II,  respectively.  Mercury  likewise  forms  two  ions, 
Hg+  and  Kg"1"*",  which  are  found  in  solutions  of  mercurous  and  mercuric 
nitrate,  HgNO3  and  Hg(NO3)2. 

The  stannous  ion  has  a  strong  tendency  to  take  on  more  positive 
charges  of  electricity,  thereby  passing  into  the  form  of  the  stannic  ion, 
while  the  mercuric  ion  can  readily  give  up  one  positive  charge,  thereby 
becoming  a  mercurous  ion ;  and  the  latter,  again,  can  give  up  its  charge 
and  become  a  neutral  mercury  atom.  When,  therefore,  a  solution  of 


54  Electrolytic  Dissociation  Thtory 

mercuric  chloride  is  treated  with  stannous  chloride  one  of  two  reactions 
will  take  place,  according  to  the  amount  of  the  latter  reagent  which  is 
used;  the  Hg^-ions  will  be  reduced  either  to  Hg+-ions,  which  unite 
with  Cl~-ions  to  give  a  white  precipitate  of  mercurous  chloride,  or  they 
will  be  completely  reduced  to  metallic  mercury,  which  appears  in  the 
form  of  a  black  precipitate  :  — 

2HgCl2  +  SnCl2  =  2HgCl  +  SnCl4  ; 

II  II  0  IV 

HgCla   +  SnCla  =  Hg        +  SnCl4. 

The  valence  of  any  element  when  in  the  uncombined  state  is  to  be 
regarded  as  zero,  as  is  indicated  for  the  metallic  mercury  appearing  in 
the  last  reaction.  The  ionic  changes  which  occur  in  these  reactions 


are:- 

++ 


Hg 

63.  Sulphur  can  exist  both  uncombined  and  in  three  different 
states  of  oxidation,  corresponding  to  the  compounds  H2S,  SO2,and  SO3, 
in  which  it  has  respectively  a  negative  valence  of  II  and  a  positive  va- 
lence of  IV  and  VI  ;  that  is,  it  has  a  valence  of  II  as  an  element  which 
is  negative  toward  hydrogen,  and  a  valence  of  IV  or  VI  as  an  element 
which  is  positive  with  respect  to  oxygen.  Corresponding  to  these 
valences,  it  forms  the  ions  S~  ,  SO3~~,  and  SO4~",  which  occur  in 
solutions  of  sodium  sulphide,  sodium  sulphite,  and  sodium  sulphate, 
respectively. 

As  has  been  seen,  the  solution  pressure  of  sulphur,  by  virtue  of 
which  it  would  pass  into  solution  as  negative  ions,  is  very  small.  There- 
fore its  ions  can  readily  be  forced  out  of  solution  to  produce  the  free 
element,  and  any  oxidizing  agent,  as,  for  instance,  chlorine  water  or 
potassium  permanganate,  will  cause  a  precipitation  of  free  sulphur  if  it 
is  added  to  a  solution  of  hydrogen  sulphide  :  — 

H2S  +  C12  =  2HC1  +  S  ; 

5H2S  +  2KMnO4  7+-  3Kf2SQ4  =  2MnSO4  +  K2SO4  +  8H2O  +  5s, 
or  S-"  +  Cl2  =  2Cl-+S; 

ioH+  +  5S-  +  56  =  5H20  +  58. 

The  free  sulphur  so  formed  is  not  easily  oxidized  further,  although  this 
can   be   brought   about,  as,  for  example,  by  boiling  with  concentrated 


I 
Oxidation  and  Reduction  55 

nitric  acid,  whereby  sulphuric  acid  is  produced.  Sulphur  can  most 
readily  be  oxidized  by  allowing  it  to  burn  in  the  oxygen  of  the  air,. 
forming  sulphur  dioxide,  and  this,  when  dissolved  in  water,  gives  sul- 
phurous acid,  which  yields  SO3~  ~-ions,  SO2  +  H2O  ^  H2SO3  ^  2H+  + 
SO3~  ~.  Sulphur  in  this  state  of  oxidation  possesses  a  valence  of  IV. 
This  fact  is  readily  seen  from  the  compound  SO2,  but  in  the  ion  SO3~" 
it  is  not  at  first  glance  so  apparent.  Three  O-atoms  possess  the  com- 
bined negative  valence  of  VI  ;  two  of  these  may,  however,  be  consid- 
ered to  be  exerted  in  holding  the  negative  charges  of  the  ion,  so  that 
four  only  remain  to  bind  the  four  positive  valences  upon  the  S-atom. 
Tetravalent  sulphur  has,  however,  quite  a  strong  tendency  to  pass 
into  some  other  state  of  oxidation,  so  that  sulphurous  acid,  if  treated 
with  any  oxidizing  agent,  is  converted  into  sulphuric  acid,  H2SO3  -f- 
H20  +  I2  —  H2S04  +  2HI,  or  SO,-"  +  O"  .+  (!,)  =  SO4-"  +  2\~. 
64.  Chromium  most  frequently  occurs  in  its  compounds  in  one  of 
two  states  of  oxidation,  represented  by  the  oxides  CrO3  and  Cr2O3,  and 
in  these  compounds  its  valence  is  either  VI  or  III.  The  compounds 
derived  from  CrO3  are  yellow  or  red,  those  from  Cr2O3  are  green  or 
blue  ;  and  whenever  a  change  in  the  state  of  oxidation  occurs  it  is  ac- 
companied by  a  change  in  color.  Potassium  bichromate,  K2Cr2O7,  is 
typical  of  the  compounds  in  which  chromium  has  a  valence  of  VI. 
Its  ions,  when  it  is  dissolved,  are  2K+  and  Cr2O7~~.  In  the  latter  the 
combined  negative  valences  of  the  oxygen  atoms  are  fourteen.  Two 
of  these  are  filled  by  the  negative  charges  on  the  ion,  thus  leaving 
twelve  to  hold  twelve  positive  valences  of  the  two  hexavalent  Cr-atoms. 
When  potassium  bichromate  is  subjected  to  the  action  of  a  reducing 
agent,  as,  for  example,  when  nascent  hydrogen  is  produced  in,  or  when 
hydrogen  sulphide  is  passed  into,  a  solution  of  it  which  contains  also* 
a  free  acid,  a  change  of  color  takes  place  from  yellow  to  green,  which 
indicates  a  reduction  :  — 

vi  in  > 

K2Cr2O7  +  8HC1  +  6H  =  2KC1  +  2CrCl3  +  ;H2O, 

or 


Cr2O7-~  +  8H+  +  6(H)  =  2Cr+  +  +  +  ;H2O. 

By  this  reaction  the  ion  Cr2O7~~,  on  conversion  into  two  Cr+  +  """-ions, 
has  lost  fourteen  equivalents  in  the  seven  atoms  of  negative  oxygen, 


56  Electrolytic  Dissociation   Theory 

but  has  lost  two  negative  charges  and  gained  six  positive  charges  of 
electricity  —  an  algebraic  total  of  eight  charges,  so  that  the  total  change 
has  been  what  may  be  called  a  loss  of  six  oxidation  equivalents.  These 
six  equivalents  of  oxygen  have  been  transferred  to  the  six  neutral  atoms 
of  hydrogen,  whereby  they  are  converted  into  water  molecules.  Thus 
in  this  as  in  all  reactions  there  is,  for  every  equivalent  of  reduction,  a 
corresponding  amount  of  oxidation.  It  is  to  be  noted  that  the  H+-ions 
which  appear  on  the  left  side  of  the  equation,  although  they  disappear 
as  such,  do  not  suffer  either  oxidation  or  reduction,  for  they  have  merely 
become  part  of  H2O-molecules. 

65.  Among  the  various  possible  oxides  of  manganese  are  MnO  and 
Mn2O7,  in  which  it  exhibits  the  valences  of  II  and  of  VII.  The  com- 
pounds of  the  former  are  colorless,  and  of  the  latter  are  a  very  intense 
reddish  violet.  When  dissolved  in  an  acid  the  oxide  MnO  yields  the 
ion  Mn++,  in  which  the  metal  has  the  same  valence,  viz.,  II:  — 

MnO  +  2H+  =  Mn++  +  H2O. 

A  characteristic  compound  derived  from  the  oxide  Mn9O7  is  potas- 
sium permanganate,  KMnO4,  which  when  dissolved  in  water  yields 
ions  as  follows  :  KMnO4  ^  K+  +  MnO4~.  In  the  ion  MnO4~  the 
metal  possesses  the  same  valence  as  in  the  oxide  from  which  it  is 
derived ;  for  of  the  eight  negative  valences  of  the  four  oxygen  atoms 
of  the  ion,  one  is  bound  by  the  negative  charge,  leaving  seven  to 
engage  the  seven  positive  valences  of  the  manganese.  If  an  acid 
solution  of  potassium  permanganate  is  reduced  by  means  of  nascent 
hydrogen  a  disappearance  of  the  deep  red  color  takes  place :  - 

2KMnO4  +  icH  +  3H2SO4  ==  K2SO4  +  2MnSO4  +  8H2O, 
or 

MnO4-  +  s(H)  +  3H+  =  Mn++  +  4H2O. 

In  this  case  it  is  seen  that  the  MnO4~-ion  is  reduced  by  five  equivalents 
on  conversion  into  Mn++,  while  five  neutral  H-atoms  have  been 
correspondingly  oxidized  with  the  formation  of  water.  Although  the 
H+-ions  have  also  combined  with  oxygen  to  form  undissociated  H2O, 
yet  they  have  thereby  suffered  no  change  in  their  state  of  oxidation. 


CHAPTER   V 

THE    MORE    COMMON    IONS    AND    THEIR    CHARACTERISTICS 

IN  order  to  assist  the  reader  in  the  specific  application  of  the 
principles  laid  down  in  the  foregoing  pages  to  the  chemical  processes 
involved  in  an  ordinary  course  of  qualitative  analysis  or  of  inorganic 
preparations,  there  is  given  in  this  chapter  a  list  of  the  most  impor- 
tant ions  which  are  formed  from  the  elements  of  common  occurrence, 
together  with  equations  showing  their  important  reactions  with  other 
ions  and  their  behavior  toward  oxidizing  or  reducing  agents.  In  most 
of  the  equations  only  those  bodies  are  included  which  actually  enter 
into  reaction,  and  the  reader  must  constantly  keep  in  mind,  when  inter- 
preting these  equations,  the  fact  that  the  electrical  charges  upon  one 
ion,  or  set  of  ions,  must  at  all  times  be  balanced  by  those  on  some 
other  ions  of  opposite  electrical  charge,  and  that  these  other  ions  do 
not  always  appear  in  the  equations.  For  example,  the  ionic  reaction 

Ag+  +  Cr  =  AgCl 

expresses  only  the  important  ionic  change  upon  which  reactions  such 
as  the  following  depend  :  — 


Ag2SO4(soln.)  +  2KCl(soln.)  =  2AgCl(ppt.)  +  K2SO4(soln.) 
AgNO3(soln.)  +  HCl(soln.)    =  AgCl(ppt.)    +  KNO3(soln.) 

It  is  evident,  then,  that  this  chapter  can  only  be  successfully 
studied  in  connection  with  the  reading  of  the  descriptive  chemistry 
of  the  elements  involved,  and  with  a  knowledge  of  the  complete 
reactions  of  which  these  are  a  simplified  form. 

No  attempt  has  been  made  to  include  in  this  chapter  detailed 
descriptions  of  the  qualitative  tests  for  ions,  but  rather  to  explain 
from  the  standpoint  of  the  Electrolytic  Dissociation  Theory  certain 
of  the  tests  which  are  described  in  any  of  the  standard  manuals  of 
qualitative  analysis. 

57 


58  Electrolytic  Dissociation   Theory 

THE  IONS  OF  THE  METALS 
Silver 

(1)  Ag+   colorless,  exists  in  acid  or  neutral  soln. 

with  Cl~ :  a  white,  curdy  ppt,  AgCl,  which  turns  black  on  ex- 
posure to  the  light.  It  is  sol.  in  NH4OH  and  KCN,  due 
to  the  formation  of  complex  ions  (2)  and  (3). 

with  S~~:   a  black  ppt.,  Ag.2S,  insol.  in  acids  or  bases. 

with  OH~:  the  fairly  strong  base  AgOH  is  formed,  which,  how- 
ever, is  but  slightly  sol.  It  separates  from  soln.  in  the  form 
of  the  brown  ppt.  Ag.2O.  This  ppt.  will  dissolve  in  pure 
water  to  such  an  extent  that  its  soln.  will  color  litmus  blue. 

(2)  Ag(NH3)2+   stable  in  an  ammoniacal  soln.     It  does  not  form  a  ppt. 

with  Cl~  nor  with  Br~,  if  a  large  excess  of  NH4OH  is  present ; 
but  with  I~,  a  yellow  ppt.,  Agl,  because  the  complex  ion  is 
still  dissociated  sufficiently  to  give  a  concentration  of  Ag+-ions 
which  with  I~-ions  will  exceed  the  solubility  product  of  Agl. 
The  solubility  product  decreases  for  the  three  salts  in  the 
order  AgCl,  AgBr,  Agl.  The  complex  ion  is  destroyed  by 
acids :  —  Ag(NH3)2+  +  2H+  =  Ag+  -f  2NH4+. 
with  S~~ :  a  ppt.  of  Ag.2S. 

(3)  Ag(CN).2~   a  very  stable  ion  in  neutral  or  alkaline  soin. 

with  C7~,  Br~,  I~:   no  ppt. 

with  S~~:  black  ppt.,  because  the  complex  ion  gives  sufficient 

Ag+-ions  to  exceed  the  solubility  product  of  Ag.2S. 
with  acids :   the  complex  ion  is  destroyed :  — 

Ag(CN)2~  -f-  H+  =  AgCN  +  HCN. 

Mercury  (ous) 

(4)  Hg+   colorless,  exists  in  slightly  or  strongly  acid  soln. 

with  Cl~~:   a  white  ppt.,  HgCl,  insol.  in  acids. 
with  S":   black  ppt.,  Hg.2S. 

Mercury  (ic) 

(5)  Hg"1"1"   colorless,  exists  in  neutral  or  acid  soln. 

with  S~~:   black  ppt.,  HgS,  insol.  in  acids  and  in  HNOg. 
with  Sn+  +  :   reduced  to  Hg+  or  to  Hg  (ppt.). 
-f  Sn++—  2Hg+-j- 


Ions  and  Their  Characteristics^  59 

Lead 

(6)  Pb"1"1"   colorless,  exists  in  neutral  or  acid  soln.  ;,  . 

with  Cl~:   a  white  ppt,  PbCl.2,  which,  however,  is  somewhat  sol. 

in  hot  water. 

with  S~~:   black  ppt,  PbS,  insol.  in  dil.  acids. 
with  OH~:    white  ppt,  Pb(OH)2,  sol.  in   an  excess  of  OH~  to 

form  the  plumbite  ion  :—  PbO2H.2  +  2OH~  =  PbO.2"~  + 

2  H2O.      PbO2H2,  like   A1O3H3,  possesses   both   basic   and 

acidic  properties  (see  A1O3H3). 

Bismuth 

(7)  Bi"1"*"-*   colorless,  exists  in  neutral  or  acid  soln. 

with  S~~:   black  ppt,  Bi2S3. 

in  dil.  and  very  faintly  acid  soln.  containing  Cl~;   a  white  ppt., 
BiOCl,  which  is  sol.  in  stronger  acid :  — 

BiOCl  +  2H+  ^  Bi+  +  +  +  Cl-  +  H2O. 

££*      ••••%'•.  - 

Copper 

(8)  Cu++   pale  blue*  exists  in  neutral  or  acid  soln. 

with  S~~:    a  dark  brown  ppt.,  CuS,  insol.  in  dil.  acid. 

with  QH~:   a  light  blue  ppt,  Cu(OH)2,  sol.  in  NH4OH,  forming 

complex  ion  (9). 
with  metallic  Cu  or  with  CN~  or  I~:    it  is  reduced  to  Cu"1" :  — 

Cu++  +  Cu  =  2Cu+ 

(CN)2  (gas) ;  2Cu++ 
with  Fe(CN)^~  ~     ~  in  acetic  acid  soln. :   a  reddish  brown  ppt. 

(9)  Cu(NH3)4"l">   intense,  deep  blue,  stable  in  presence  of  slight  excess 

of  NH4OH. 

with  S~  ~:   brown  ppt.,  CuS,  since  complex  ion  is  somewhat  dis- 
sociated into  simple  Cu^^-ions. 

with  CN~  in  alkaline  soln. :   it  is  reduced  with  the  formation  of 
complex  ion  (10)  while  the  blue  color  disappears:  — 
2Cu(NHs)4++  +  6CN-  —  2Cu(CN)2-  +  (CN)a  +  8NH3. 

(10)  Cu(CN).2~   colorless,  stable  in  alkaline  soln. 

with  S  ~  ~:   no  ppt. 

In    this    ion    (which    is    analogous    to    Ag(CN)2~),    copper    is 
monovalent,  as  in  the  oxide  C.uaO. 


60  Electrolytic  Dissociation  Theory 

Cadmium 

(n)     Cd++   colorless,  exists  in  neutral  or  acid  soln. 

with  S~~:   a  yellow  ppt.,  CdS,  insol.  in  dil.  acids,  in  NH4OH, 

or  in  KCN,  but  sol.  in  strong  acids. 

with  OH~:   a  white  ppt,  Cd(OH)2,  sol.  in  NH4OH  with  formation 
of  (12). 

(12)  Cd(NH8)4"l"+   colorless,  exists  in  ammoniacal  soln. ;  it  is  so  much  dis- 

sociated into  simple  ions  that  a  ppt.  is  formed  with  S~~;  it 
is  not  reduced  by  CN~  or  I". 

(13)  Cd(CN)4~~   colorless,   exists  in   alkaline   solns. ;   it    is   considerably 

dissociated  into  simple  Cd++-ions,  so  that  a  ppt.  is  formed 
with  S"". 


Arsenic  exists  in  different  states  of  oxidation.  Its  trioxide,  As2O3,  shows 
weakly  basic  properties  in  that,  with  cone,  strong  acids,  the 
As+  +  +-ion  is  produced.  Its  acidic  properties  are  much  more 
pronounced  in  that  it  dissolves  in  water  to  form  arsenious  acid, 
H8AsO3,  and  in  bases  to  form  salts  of  that  acid.  The  pentoxide 
shows  almost  exclusively  acidic  properties  in  that  it  will  not,  on 
dissolving,  yield  positive  ions,  but  will  form  only  arsenic  acid, 
H3AsO4,  or  its  salts. 

(14)     As+  +  +   colorless,  exists  in  strongly  acid  soln. 

with  S~  ~:  a  yellow  ppt.,  As2S3,  insol.  in  acids,  but  sol.  in  alkalies, 
or  alkaline  sulphides  :—  As2S3  -)-  38"  ~  =  2AsS3         ; 
As2S3  +  3Sx--=2AsS4---+(3X-  5)s. 

with  OH~:   forms  arsenious  acid,  AsO3H3,  soluble,  which  with 
more  OH"  gives  the  arsenite  ion:-  — 

As03H3  +  3OH-  =  As03--      4-3H20. 


(15)  AsO3~~~    exists  in  neutral  or  alkaline  soln. 

with  Jf+:  undissociated  AsO3H3  is  formed,  which,  with  a  large 
excess  of  H+-ions,  reacts  as  a  base  with  a  partial  formation 
of  As+  +  +-ions:—  As(OH)3  +  3H+  ^  As+  +  +  +  3H2O. 

(16)  AsO4~"  "   exists  in  neutral  or  alkaline  soln. 

with  Mg+  +  and  NH  +  in  strongly  ammoniacal  soln.:  a  white 
crystalline  ppt.,  MgNH4AsO4. 


Ions  and  Their  Characteristics  61 

with  S~~  in  acid  soln.  :  is  reduced,  very  slowly  in  the  cold,  more 
rapidly  when  hot  :  —  AsO4  ---  +  S~  ~  +  8H+  =  As+  +  +  -f 
4H2O  -|-  S  ;  in  a  cold,  strongly  acid  soln.  a  ppt.  forms 
slowly,  which  is  almost  exclusively  As2S5:  —  2AsO4~~~-J- 
58-"  +  i6H+  =  As2S5  +  8H2O. 

Antimony  is  very  similar  to  arsenic  as  regards  the  compounds  which  it  forms 
and  their  characteristics.  It  is,  however,  distinctly  more  metallic 
in  character.  Its  trioxide  is  principally  basic  in  its  nature, 
although  its  pentoxide,  like  that  of  arsenic,  is  acidic. 

(17)  Sb"*"*"4"   colorless,  exists  in  acid  soln. 

with  S~~:   bright  orange  red  ppt.,  Sb2S3,  which  is  sol.  in  alkalies 
or  alkaline  sulphides  :  —  Sb2S3  -f-  38"  ~  =  2SbS3         ; 
SbaS,  +  3SX-  -  =  2SbS4-  -  -  +  (3x  -  5)8. 

with  Cl~  in  faintly  acid  soln.  :  a  white  ppt.  of  antimony  oxy- 
chloride,  SbOCl,  which,  however,  is  sol.  in  more  cone,  acid  :  — 

Sboci  +  2H+  ^  sb+++  +  cr  +  H2o. 

with  OH~:  white  ppt.,  Sb(OH)3,  insol.  in  excess  of  ammonia, 
and  sol.  only  in  cone,  caustic  alkalies. 

Tin 

(18)  Sn"1"*   colorless,  exists  in  faintly  or  strongly  acid  soln. 

with  S~~:  dark  brown  ppt.,  SnS,  insol.  in  dil.  acids,  alkalies, 
and  alkaline  sulphides,  but  sol.  in  presence  of  Sx~  "-ions  in 
consequence  of  an  oxidation:  — 

SnS  +  Sx-  -  =  SnS8-  "  +  (x  -  2)8. 

with  oxidizing  agents  :   changes  with  great  readiness  to  Sn+  +  +  +  :  — 


Sn++  -f  °  +  2H+  =  Sn+  +  +  +  +  H2O. 

with   OH~:  white  ppt.,   Sn(OH)2,   sol.   in   an  excess   of   strong 
alkali  :  —  SnO2H2  +  2QH-  —  SnO2~  "  +  2H2O. 

(19)     Sn"1"1"1"*"   colorless,  exists  in  acid  soln. 

with  S~  ~:  yellow  ppt.,  SnS2,  insol.  in  dil.  acids,  but  sol.  in  alkaline 

sulphides  :  —  SnS2  +  S~  ~  =  SnS8~  ~. 
with  OH~:  white  ppt.,  Sn(OH)4,  sol.  in  excess  of  alkali  :  — 

Sn(OH)4  +  2OH~  =  SnO3~-  +  3H2O. 
with  metallic  tin  :   it  is  reduced  :  —  Sn+  +  +  +  -|-  Sn  =  2Sn+  +. 
with  Zn  :    Sn+  +  +  +  is  reduced  to  Sn+  +,  and  then,  in  a  soln.  not 
more  than  faintly  acid,  to  metallic  Sn.     Sn,  however,  dis- 
solves in  acids  :  —  Sn  -|-  2H+  =  Sn++  -|-  H2- 


62  Electrolytic  Dissociation   Theory 

Irbn 

(20)     Fe++     faint  green,  exists  in  neutral  or  acid  soln.  ;  it  is  easily  oxidized 


with  S~  ~  in  alkaline  soln.  :   black  ppt..  FeS,  sol.  in  acids. 

with  OH~:  greenish  ppt.,  Fe(OH)2. 

with  Fe(CN\          .-    white   ppt.,    Fe2[Fe(CN)6],   which   rapidly 

turns  blue  in  consequence  of  oxidation  from  the  oxygen  of 

the  air. 

with  Fe(  CW)6  ---  .-  deep  blue  ppt.  of  Turnbull's  blue,Fe3[Fe(CN)6]2. 
with  SCJV~:  no  reaction. 
with  COS~  ~:  white  ppt.,  FeCO3,  insol.  in  neutral  or  alkaline  soln. 

(21)     Fe+"f"*"    colorless,1  exists  in  faintly  or  strongly  acid  soln. 
with  S~~  in  acid  soln.  :  is  reduced  :  — 

2Fe+  +  ++  S-~=  2Fe++  +  S. 

with  S~~  in  alkaline  soln.:  black  ppt.,  FeS,  with  a  simultaneous 
production  of  free  sulphur:  — 

2Fe+  +  +  +  38--  =  2FeS  -f  S. 
with  OH~:    red  ppt.,  Fe(OH)8. 

with  Ac~  in  acetic  acid  soln.  when  boiled  :  a  red  ppt.  of  basic  ferric 
acetate:- 


Fe+  +  +  +  Ac'  +  2OH-  (from  H2O)  =  Fe-OH 

NAc. 

with  PO4        .•  a  white  ppt.,  FePO4,  insol.  in  alkalies  or  in  acetic 
acid,  but  sol.  in  strong  acids  :  — 

FePO4  +  H+  =  Fe+  +  +  +  HPO4~  ~. 

•with  jR?(CW)6        /    a  sol.  olive  green  complex  compound. 

with  Fe(CN)s  ----  /  a  blue  ppt.  of  Prussian  blue,  Fe4[Fe(CN)6]3. 

with  SCN~  in  acid  soln.  :    a  sol.  intensely  blood-red,  non-ionized 
compound,  Fe(SCN)8. 

with  solid  suspended  BaCO*:  a  ppt.  of  Fe(OH)8,  which  may  be 
accounted  for  as  follows:  the  ferric  ion  hydrolyzes  easily 
with  water,  Fe+  +  +  -f  3H2O  ^  Fe(OH)3  +  3H+  (see 
Hydrolysis,  pp.  42-44)  ;  but  the  H+-ions  produced  by  this 
reaction  cannot,  in  the  presence  of  suspended  BaCO3,  attain 
any  appreciable  concentration,  since  they  react  with  the 
latter  to  form  CO2  and  H2O  (see  Solubility  of  Carbonates 
in  Acids,  p.  39).  Hence  the  hydrolysis  of  a  ferric  salt  which 
,in  pure  water  reaches  an  equilibrium  before  any  ppt.  of 

1  Neutral  solutions  of  ferric  salts  are  usually  brown,  due  to  the  presence  of  small 
amounts  of  Fe(OH)3  resulting  from  hydrolysis.  The  yellow  color  of  ferric  chloride  solu- 
tions is  due  to  the  presence  of  undissociated  FeCl3. 


Ions'  and  TJieir  Characteristics  63 

Fe(OH)3  is  produced,  can,  in  the  presence  of  BaCO3,  con- 
tinue to  completion,  and  thus  a  practically  complete  pre- 
cipitation of  ferric  iron  can  occur. 

(22)  Fe(CN)6 a  very  stable  ion  which  gives  ppts.  with  most  of  the 

simple  ions  of  the  heavy  metals. 

with  oxidizing  agents  :   it  is  readily  converted  into  the  ferricyanide 
ion  :—  2Fe(CN)6~-      '  +  C1a  =  2Fe(CN)6"  "    '  +  2Cr. 

(23)  Fe(CN)6~"  "   a  very  stable  ion  which,  similarly  to  the  ferrocyanide 

ion,  forms  ppts.  with  the  simple  ions  of  most  of  the  heavy 
metals. 

Aluminum 

(24)  Al"l"h+   colorless,  exists  in  faintly  or  strongly  acid  soln. 

with  OH~:   a  flocculent,  white  ppt.,  A1(OH)3,  insol.  in  ammonia 

but  sol.  in  strong  alkalies. 

A1O3H3  has  both  weakly  acidic  and  weakly  basic  properties ;  that  is, 
the  minute  quantity  which  dissolves  in  pure  water  dissociates  in 
two  different  ways,  whereby  H+-  and  OH~-ions,  respectively,  are 
produced :  — 
(i)     Al(OH),  ^  Al*  +  +  +  3OH- 

j  A103H8    ^  3H+      +  A10,-  -  or 

V  (  A1(OH),  ^  HA1O2  +  H2O  ;  HA1O2  ^  H+  -f  A1O~2. 
Hence  its  solubility  in  acids  and  bases ;  for  in  the  former  case 
the  H+-ions  of  the  acid  destroy  the  OH~-ions  produced  by  the 
dissociation  of  the  aluminum  hydroxide,  thus  allowing  the  disso- 
ciation (also  the  solution)  to  proceed  to  completion ;  while  in  the 
second  case  the  OH~-ions  of  the  base  destroy  the  H+-ions  pro- 
duced   by   the    dissociation    of    the   aluminic    acid    (H8A1O8   or 
HA1O2),  thus  allowing  the  latter  to  continue  to  completion :  — 
(i)     A1(OH)3  +  3H+     =  A1+  +  +      +  3H20. 
H3A103    +3OH-  =  A103--   -  +  3H2Oor 
Al(OH),  +  OH'    =  A1O2-       +  H2O. 

Ammonium  hydroxide,  however,  does  not  yield  OH~-ions  in  suffi- 
cient concentration  to  react  with  the  very  few  H+-ions  from  the 
H3A1O3,  which  accounts  for  the  failure  of  the  latter  to  readily 
dissolve  in  that  reagent. 
with  Ac~  in  very  faintly  acid  soln.  when  boiled :  a  flocculent,  white 

ppt.,  Al(OH)2Ac. 
with  BaCOz:  ppt.  of  A1(OH)3  (see  action  of  BaCO3  on  Fe 


64  Electrolytic  Dissociation   Theory 

Chromium 

(25)  Cr++   can  exist  in  neutral  or  acid  soln.,  but  it  changes  with  great 

readiness  to  Cr+  H  +  in  consequence  of  absorption  of  oxygen 
from  the  air. 

(26)  Cr+H   h   violet  or  green  in  color,  exists  in  faintly  or  strongly  acid  soln. 

with  OH~:  a  flocculent,  light  green  ppt.,  Cr(OH)3,  insol.  in 
ammonia  and  alkalies  when  hot,  but  sol.  when  cold  in 
excess  of  alkali.  CrO3H3,  like  A1O3H3,  possesses  both 
basic  and  acidic  properties  (see  A1O8H8). 

with  Na-iOi  and  other  oxidizing  agents  in  alkaline  soln. :  is  oxidized 
to  the  yellow  chromate  ion  :  — 

2Cr+  +  ++  ioOH-  +  3O  =  2CrO4-~  +  5H2O. 

with  BaCOz:   ppt.  of  Cr(OH)3  (see  action  of  BaCO3  on  Fe+  +  +). 

(27)  CrO4~~   in  which  Cr  possesses  the  valence  of  VI ;  yellow;  exists  in 

neutral  or  alkaline  soln. 
with  Ba+  +  and  Pb+  +  in  neutral  or,  in  acetic  acid  soln. :    yellow 

ppts.,  BaCrO4,  PbCrO4. 

with  Ag*  in  neutral  soln. :  red  ppt.,  Ag2CrO4. 
with  Jf+:   it  is  converted  partially  or  wholly  into  the  bichromate 

ion  :  —  2CrO4-  ~  -f  2H+  ^  Cr2(V  "  +  H2O. 

(28)  Cr2O7~~"   in  which,  as  in  CrO4~~,  Cr  possesses  the  valence  of  VI ; 

red  when  concentrated;  exists  in  neutral  or  acid  soln. 
with  OH~:   changes  to  the  chromate  ion  :  — 

Cr2O7-"  +  2OH~  =  2CrO4~-  +  H2O. 

with  S~~  and  with  other  reducing  agents  in  acid  soln. :  is  reduced 
to  the  chromic  ion  (see  p.  65)  :  — 

+  +  7H2O  +  3S. 


Cobalt 

(29)     Co++   pink,  exists  in  neutral  or  acid  soln. 

with  S~~:   a  black  ppt.,  CoS,  insol.   in  alkalies  or>   after  once 

formed,  in  dil.  acids. 

with  OH~:  a  light  blue  ppt.,  Co(OH)2,  insol.  in  excess  of 
ammonia  or  strong  alkalies,  but  sol.  in  NH4C1  (probably 
due  to  the  great  decrease  of  the  number  of  OH~-ions  in 
consequence  of  the  NH4+-ions  from  the  NH4C1). 


t 

Ions  and  Their  Characteristics  65 

with  CN~  in  neutral  soln.  :  a  reddish  brown  ppt.,  Co(CN)2r 
which  with  excess  of  CN~  redissolves  to  form  the  complex 
cobalto-cyanide  ion,  Co(CN)6  (sol.),  which  by  means  of 

oxidizing  agents  is  readily  converted  into  the  cobalti-cyanide 
ion,  in  which  Co  has  the  valence  III :  — 

2Co(CN)6--      "  +  Cla  =  2Co(CN)6--  '  +  2C1-. 

(30)  Co(CN)6          an  ion  which  can  exist  in  both  acid  or  basic  solns.     It 

is  comparable  in  stability  with  the  SO4~~-ion. 
with     Cu+  "*"    in    acid    soln. :    a    very    insol.     light     blue     ppt., 
Cu3[Co(CN)J2. 

Nickel 

(31)  Ni++   green,  exists  in  acid  or  neutral  soln. 

with  S~~:   a  black  ppt.,  NiS,   insol.   in    alkalies  or,   after  once 

formed,  in  dil.  acids. 
with   OH~:   a    light   green    ppt.,   Ni(OH)2,  insol.    in   excess   of 

ammonia    or    strong    alkalies,    but,    like    Co(OH)2,    sol.    in 

NH4C1. 
with  CN~  in  neutral  soln. :    a  green  ppt.,  Ni(CN)2,  which  with 

excess  of  CN~  redissolves  to  form  the  complex  nickelo-cyanide 

ion,  Ni(CN)4~~  (sol.),  which  is  stable  in  alkaline  or  neutral 

soln.,  but  with  oxidizing  agents  yields  a  black  ppt.  of  nickelic 

hydroxide :  — 

Ni(CN)4-  -  +  sci  +  3OH-  =  Ni(OH)3  +  2(CN)2  +  5cr. 

Manganese 

(32)  Ma**   colorless,  exists  in  neutral  or  acid  soln. 

with  S~~:  a  flesh-colored  ppt.,  insol.  in  water  and  in  alkalies,  but 

sol.  in  dil.  acids  and  in  acetic  acid. 
with  OH~:  a  white  ppt.,  Mn(OH)2,  insol.  in  excess  of  alkali,  but, 

like  Co(OH)2,  sol.  in  presence  of  NH4-salts. 
with  oxidizing  agents  in  alkaline  soln.  :   a  black  ppt.  of  MnO2 :  — 

Mn++  +  Br2  +  4OH-  =  MnO2  +  2H2O  +  2Br~. 
with  PbO%  in  HNO%  soln. :   a  red  color  of  permanganate  ion  :  — 

2Mn++  +  5PbO2  +  4H+  =  2MnOr  +  5?b++  +  2H2O. 
with  CN~:   a  brownish  ppt.,  which  with  a  large  excess  of  CN~-ions 

redissolves  to  a  slight  extent  with  formation  of  the  unstable 

ion,  Mn(CN)6 . 

with  COi~~:   a  white  ppt.,  MnCO3,  insol.  in  NH4-salts. 

with  PO±          and  NH*  in   ammoniacal  soln. :   a   pink  ppt.  of 

MnNH4PO4. 


66  Electrolytic  Dissociation   Theory 

(33)  Mn+  *  +  a  very  unstable  ion. 

(34)  MnO4~"   in  which  Mn  has  a  valence  of  VI,  stable  in  alkaline  solns., 

to  which  it  imparts  an  intense  green  color. 

with  H~-ions :   it  is  unstable,  and  is  in  part  oxidized  to  MnO4~ 
at  the  expense  of  the  other  part,  which  is  reduced  to  MnO2 :  — 
3MnO4—  +  4H+  =  2MnO4~  +  MnO2  +  2H2O. 

(35)  MnO4~   in  which  Mn  has  a  valence  of  VII,  stable  in  neutral  or  acid 

soln. ;  has  an  intense  reddish  violet  color. 
in  alkaline  soln.  when  boiled:  it  changes  to  the  green  MnO4~~- 

ion  :  —  2MnO4~  +  2OH~  =  2MnO4~  "  -f-  H2O  -f  O. 
Both  MnO4~  ~-  and  MnO4~-ions  are  strong  oxidizing  agents.     In 

alkaline  soln.  they  are  reduced  to  MnO2,  a  black  ppt. ;  in 

acid  soln.  to  Mn+  "*" :  — 

MnO4~  +  5Fe++  +  8H+  =  Mn++-f-  5Fe+-f+  +  4H2O. 

Zinc 

(36)  Zn++   colorless,  exists  in  neutral  or  acid  soln. 

with  S~~:   a  white  ppt.,  ZnS,  insol.  in  alkalies,  in  neutral  soln., 
and  in  acetic  acid,  but  sol.  in  strong  acids. 

with    OH~:   a  white  ppt.,  Zn(OH)2,  which   is  sol.   in    acids,  in 
strong  alkalies,  and   in  ammonia.      ZnO2H2  possesses  both 
acidic  and  basic  properties  (see  A1O3H3).     It  dissolves  in 
strong  alkalies  with  the  formation  of  the  negative  zincate 
ion:—  ZnO2H2  +  2OH~  =  2H2O  -f  ZnO-f. 
Ammonium  hydroxide  is,  however,  not  sufficiently  dissociated 
into  OH~  to  cause  the  above  reaction  to  take  place ;  the 
solubility  in   ammonia   is   due  rather  to   the   formation   of 
the  complex  zinc-ammonia  ion  :  — 
Zn(OH)2  -f  4NH4OH  =  Zn(NH3)4++  -f  2OR-  +  4H2O. 

Calcium 

(37)  Ca++   colorless,  exists  in  neutral,  acid,  or  weakly  alkaline  soln. 

with  OH~  when  concentrated :    white  ppt.,  Ca(OH)2. 

with  COi~  ~:   white  ppt,  CaCO3,  which,  however,  is  sol.  in  very 

weak  acids:—  CaCO3  +  H+  =  Ca++  +  HCO3~. 
with  CZO^~~:  a  white  ppt.  of  calcium  oxalate,  CaC2O4,  sol.  in 

strong  acids. 
with  SO*~~:   a  white  ppt.,  CaSO4,  which  is  not  entirely  insol.  in 

water. 


Ions  and  Their  Characteristics  6y 

Strontium 

(38)  Sr"1"*"   colorless,  exists  in  neutral,  acid,  or  alkaline  soln. 

with  COS~  "•    a  white  ppt.,  SrCO3. 

with  SO4~~:   a  white  ppt.,  SrSO4,  much  less  sol.  than  CaSO4. 
Barium 

(39)  Ba+  +   colorless,  exists  in  neutral,  acid,  or  alkaline  solution. 

with  CO3~  ~:   white  ppt,  BaCO3. 

with  CrO^~  ~   in   alkaline,    neutral,    or   acetic   acid  soln. :   yellow 

ppt.,  BaCrO4. 
with   SO4~  ~  in    neutral,   acid,   or  alkaline  soln. :    white   ppt.   of 

BaSO4. 
Magnesium 

(40)  Mg+  +   colorless,  exists  in  neutral  and  acid  soln. 

with  OH~:  a  white  ppt.,  Mg(OH)2,  which  is  easily  sol.  in  NH4- 
salts,  probably  because  in  the  presence  of  NH4+  the  OH~-ions 
cannot  reach  a  sufficient  concentration  to  give  the  solubility 
product  with  Mg+  +  (see  Co(OH)2). 

with  COa~~  and  C2O4~~:   no  ppt.,  if  NH4-salts  are  also  present. 

with  NH*  and  PO± in  ammoniacal  soln. :  white  crystalline 
ppt.,  MgNH4P04. 

Sodium,  Potassium,  and  Ammonium 

(41)  Na+,  K+,  NH4+   colorless,  exist  in  neutral,  acid,  or  alkaline   soln. ; 

do  not  form  complex  ions,  do  not  form  any  insol.  ppts.  with 

the  common  anions. 
K+   gives   with    PtCl6~~    and   with    Co(NO2)6          characteristic 

difficultly  sol.  yellow  ppts.,  K2PtCl6  and  K3Co(NO2)6,  while 

the  corresponding  Na-salts  are  sol. 
NH4+  gives   with    PtCl6  likewise   a   difficultly,  sol.  yellow   ppt., 

(NH4)2PtCl6;  but  it  differs  markedly  from  K+  and  Na+  in 

that  it  forms  with  OH~  the  very  slightly  dissociated  body, 

NH4OH,  while  the  hydroxides  of  the  other  two  metals  are 

highly  ionized. 

THE  IONS  OF  THE  NON-METALS 

Boron  forms  no  simple  ions.  From  its*  trioxide,  B2O8,  is  derived  the  very 
weak  boric  acid,  B(OH)8,  which  when  treated  with  alkalies  gives 
salts  of  metaboric  acid,  HBO2 :  —  B(OH)3  ^  HBO2  +  H2O ; 
HBO2  +  OH"  ^  H2O  +  BO2~.  The  ion,  BO2~,  is  capable  of 
existence  only  in  alkaline  soln. 


63  Electrolytic  Dissociation   Theory 

Carbon   forms  no  simple  ion. 
CO3~  ~~   exists  in  alkaline  soln. 

with  ff+:  first  HCO3  is  formed,  which  with  more  H+  gives  the 
undissociated  carbonic  acid  :  —  H"1"  -\-  CO3~  ~  =  HCO8~  ; 
HC03-  4-  H+  ^  H2C03  ^  H20  +  CO,. 

with  the  ions  of  most  of  the  metals  :  ppts.  which  are  sol.  in  dil. 
acids. 

HCO3~   exists  only  in    neutral   soln.      It  does   not   form   ppts.  with   the 
ions   of   the    metals   unless   it   first   breaks    down    to    give 

co3-  - 

C2H3O2~   forms  no  insol.  salts  except  the  slightly  sol.  Ag(C.2H3O2)   and 
certain  basic  acetates. 

c2o4-- 

with  Ca++:  white  ppt.,  CaC2O4,  which  is  sol.  in  dil.  acids,  due 
to  the  formation  of  undissociated  oxalic  acid,  H2C2O4. 

Silicon  resembles  boron  in  its  tendency  to  form  ions  and  in  the  character 
of  the  ions  formed.  From  its  dioxide,  SiO2,  there  is  derived 
silicic  acid,  H2SiO8,  as  in  a  somewhat  similar  manner  carbonic 
acid,  H2CO3,  is  derived  from  CO2.  Treated  with  alkalies  it 
gives  the  silicate  ion,  which  is  capable  of  existence  only  in  alkaline 
soln.  :  —  H2SiO8  +  2OH~  ^  2H2O  +  SiO8.~  " 


Nitrogen   forms  no  simple  ions. 

NO2~   colorless,  forms   no   insol.    compounds   except    the    difficultly    sol. 
AgNO2.     In  acid  soln.  it  forms  with  H+  the  weak  nitrous 
acid,  HNO2.     In  such  a  soln.  it  will  reduce  strong  oxidizing 
agents,  whereby  it  is  converted  to  the  nitrate  ion  :  — 
2MnO4~  +  6H+  4-  5NO2~  =  2Mn++  +  $NO3~  -{-  3H2O  ; 
or  it  will  oxidize  reducing  agents,  whereby  it  is  converted 
to  nitric  oxide:  — 

2!-  4-  4H+  4-  2NO2~  =  I2  4-  2NO  4-  2H2O. 
It  is  less  stable  than  the  NO8~-ion,  and  is  hence  a  more 
active  oxidizing  agent. 

NO3~  colorless,   forms  no  insol.  compounds  ;  in  acid  soln.  it  is  a  strong 
oxidizing  agent:  — 

4-  3Fe+  +  4-  4H+  =  3Fe+  +  +  +  2H2O  4-  NO. 


Ions  and  Their  Characteristics  69 

CN~    colorless,  exists  in  slightly  alkaline  and  alkaline  soln.     It  is  excess- 
ively poisonous.     It  has   a  marked  tendency  to  form  com- 
plex ions  with  the  ions  of  the  heavy  metals.     It  forms  with 
H+  the  very  weak  hydrocyanic  acid,  HCN.     It  acts  under 
some  circumstances  as  a  reducing  agent  (see  action  on  Cu+  +). 
SCN~   colorless,  exists  in  neutral,  acid,  and  alkaline  soln. 
with  Ag+:   white  ppt,  AgSCN. 

with  J?e+  +  +:   an  intensely  red,  sol.,  but  non-ionized,  compound, 
Fe(SCN)3. 

Phosphorus  forms  no  simple  ions,  but  with  oxygen  it  forms  several,  of 
which  the  most  important  are  those  derived  from  phosphoric 
acid,  H3PO4.  This  acid  dissociates  in  three  successive  stages  :  — 

H3PO4  ^  H+  +  H2PO4- ;  H2PO4~  F^  H+  +  HPO4~  ~ ; 

HPO4-  -  ^  H+  -f  PO4 . 

In  the  first  stage  the  dissociation  is  considerable,  yielding  thus 
a  moderately  strong  acid ;  in  the  second  it  is  very  small ;  in  the 
third  it  is  excessively  small,  so  that  HPO4~~  is  a  very  weak  acid. 
PO4  ,  in  neutral  or  alkaline  soln.,  gives  ppts.  with  the  ions  of 
most  of  the  metals  (see  Mg++),  but  these  are  sol.  in  dil.  acids 
on  account  of  the  formation  of  undissociated  HPO4~~  or  H2PO4~. 

Sulphur 

S~~  exists  in  alkaline  and  neutral  soln.,  but  only  in  very  small  cone,  in 
acid  soln.,  since  in  presence  of  H+  it  forms  the  very  weakly 
ionized  H2S.  It  gives  ppts.,  mostly  deeply  colored,  with  the 
ions  of  the  heavy  metals.  It  acts  readily  as  a  reducing  agent, 
whereby  it  is  itself  oxidized  to  free  S  (see  Fe+  +  +). 
SO3~  ~  exists  in  neutral  and  alkaline  soln.,  and  to  a  certain  concentration 

in  acid  soln. 
with  J?a+  +  and  Pb*+:   white  ppts.  which  are  easily  sol.  in  dil. 

acids,  due  to  formation  of  undissociated  H2SO8. 
with  oxidizing  agents :   it  is  converted  to  SO4~  ~:  — 

so8--  +  o  —  so4-~. 

SO4~~   exists  in  acid,  neutral,  and  alkaline  soln. 

with  Pb+  +  and  with  jfta++:   white  ppts.,  PbSO4,  BaSO4,  insol.  in 

acids  and  alkalies, 
exists  in  neutral  and  alkaline  soln. 
with  H+:   it  is  decomposed  :  — 

H2S2O3 >  H2O  +  SO2  +  S. 


70  Electrolytic  Dissociation  Theory 

Fluorine 

F-  with  Ca++,  £a++,  Pb++:   ppts.  of  CaF.2,  BaF2,  PbF.2,  which,  how- 

ever, are  sol.  in  dil.  strong  acids  on  account  of  the  formation 
of  undissociated  HF. 

•'.    .    \  •      -   :':  •    '    .  ,.    X  '  ''X';.    /  •       :.•'.:,•     -  \..'  .  'J.r'   :  ' 

Chlorine 

Cl"   forms  but  few  insol.  comps.  (see  Ag+,  Hg+,  and  Pb++). 

with  strong  oxidizing  agents :  it  is  changed  to  free  chlorine,  e.g. :  — 

2Cr  +  MnO2  +  4H+  =  Mn^^  +  2H.2O  +  CL2. 
with  oxidizing  agents  in  alkaline  so  In. :   it    is   converted   to  the 
hypochlorite  ion,  C1O~,  and  to  the  chlorate  ion,  C103~:  — 
Cl-  +  Cl,  +  20H-  =  CIO'  +  2C1-  +  H20  ; 

cr  4-  3cia  +  6OH-  =  cio3-  +  6cr  +  3H2o. 

C1O8-   forms  no  insol.  comp. 

with  reducing  agents :   it  is  changed  to  the  simple  ion  :  — 
CIO,"  +  3S03-  -  =  3S04-  -  +  C1-. 

Bromine   forms  ions  analogous  to  those  of  chlorine. 

Iodine    forms  ions  analogous  to  those  of  chlorine  and  bromine ;  it  is  a  strong 
reducing  agent,  e.g.  : —  2!"  -j-  2CU"1"*"  =  2Cu+  -|~  ^ 


••f 


CHAPTER    VI 

EXPERIMENTS    ILLUSTRATIVE    OF     THE     THEORY     OF     ELECTROLYTIC 
DISSOCIATION,    AND    SOME    OF    ITS    APPLICATIONS 

THE  following  series  of  experiments  is  designed  to  accompany  the 
study  of  the  preceding  text.  The  important  points  to  be  learned  from 
each  experiment  are  indicated  by  the  italicized  questions  to  be  answered 
by  the  student.  The  actual  performance  of  the  experiments  will  be 
found  to  be  comparatively  simple,  but  satisfactory  laboratory  notes 
should  contain  much  more  than  the  mere  records  of  observations,  and 
such  notes  can  only  be  written  after  careful,  thoughtful  study  of  the 
related  portions  of  the  foregoing  chapters.  The  student  should  also 
consult  frequently  with  his  instructors,  and  explanations  and  correc- 
tions on  the  part  of  the  latter  must  be  painstaking  and  complete  if 
the  laboratory  experimentation  is  to  serve  its  purpose. 

Apparatus.  —  The  only  special  apparatus  required  for  these  experi- 
ments is  that  for  the  testing  of  electrical  conductivity.1 

Figure  5  is  a  diagram  of  the  entire  apparatus  as  mounted  on  a 
standard.  The  terminal  wires,  as  represented  by  the  symbols  +  and  - 
may  be  extended  by  means  of  a  lamp  cord  and  plug,  the  latter  of  which 
may  be  screwed  into  a  lamp  socket  of  a  i  ro-volt,  incandescent  lighting 
system.  Either  direct  or  alternating  current  may  be  used,  although  the 
former  is  somewhat  to  be  preferred. 

The  terminal  wires,  as  represented,  extend  along  the  length  of  the 
standard.  One  of  them  is  connected  with  one  .binding  screw  of  each 
of  three  incandescent  lamp  sockets ;  the  other  wire  is  attached  to  one 
binding  post  of  each  of  three  pairs.  The  other  binding  post  of  each 
pair  is  connected  with  the  other  binding  screw  of  the  lamp  socket  above 
it.  From  the  pairs  of  binding  posts  are  suspended  three  different  forms 
of  electrodes,  designated  A,  B,  and  C.  The  apparatus  is  thus  so 

1  This  form  of  apparatus  was  designed  in  the  laboratories  of  the  Massachusetts  Institute 
of  Technology,  where  it  has*  been  in  successful  use  for  two  years. 

71 


72 


Electrolytic  Dissociation   Theory 


ABC 

FIGURE  5 

arranged  that  if  a  substance  of  unknown  conductivity  is  introduced 
between  two  electrodes  belonging  to  one  of  the  pairs,  an  electric  cur- 
rent will  pass  in  proportion  to  the  conductivity  of  the  substance.  Since 
this  current  must  also  pass  through  the  lamp  directly  above,  this  will 
glow  with  a  brilliancy  roughly  proportional  to  the  amount  of  current. 

Electrodes  A  consist  of  stout  copper  wires,  and  are  to  be  used  in 
testing  the  conductivity  of  solid  substances,  lumps  of  which  are  held 
with  the  fingers  in  such  a  way  as  to  come  into  contact  with  each 
electrode,  A  i6-candle  power  lamp  may  be  used  in  the  socket  above. 

Electrodes  B  consist  of  similar  copper  wires,  but  bent  far  enough 
apart  so  that  they  will  pass  into  the  two  arms  of  a  U-tube  8  mm. 
internal  diameter  and  7  cm.  high,  when  the  latter  is  raised  from  under- 
neath. A  5<D-candle  power  lamp  (or,  somewhat  less  satisfactorily,  a 
32-candle  power  lamp)  should  be  used  with  these  electrodes,  which 
are  then  designed  to  show  differences  in  conductivity  among  good 
conductors. 

Electrodes  C  consist  of  fine  platinum  wires  supported  upon  glass 
rods1^  as  shown  in  Figure  6,  and  are  to  be  used  with  a  i6-candle  power 

1 A  pattern  of  Electrode  C  has  been  furnished  to  the  L.  E.  Knott  Apparatus  Company-, 
Boston,  who  are  in  a  position  to  furnish  this  part  of  the  apparatus.  The  remainder  of  the 
apparatus  may  be  constructed  by  any  carpenter. 


Experiments 


73 


0 

FIGURE  6 


lamp.  They  are  to  be  used  in  testing  the  conductivity 
of  liquids,  the  latter  to  be  placed  in  3-inch  test  tubes, 
which  may  be  raised  until  the  electrodes  are  immersed 
in  the  liquid.  Electrodes  A  could  be  used  in  the  same 
way,  but  have  the  disadvantage  that  they  are  strongly 
attacked  by  the  liquid  during  the  passage  of  the  current. 

1 .  The  Electrical  Conductivity  of  Different  Types 
of  Substances.  —  Test  the  conductivity  of  water,  alco- 
hol, and  pure  acetic  acid. 

Select  any  two  solid  salts  in  the  laboratory.  Test 
the  conductivity  of  each  by  holding  a  small  lump  so  as 
to  touch  both  electrodes.  Dissolve  not  more  than  0.25 
gram  of  each  in  a  3-inch  test  tube  of  pure  water,  and 
test  the  conductivity  of  the  solution. 

Test  the  conductivity  of  diluted  solutions  of  sodium 
hydroxide  and  of  hydrochloric  acid.  Dilute'  about  0.5  c.c. 
of  pure  acetic  acid  with  water  in  a  small  test  tube,  and 
test  its  conductivity. 

Dissolve  a  little  sugar  in  water ;  also  mix  about  I  c.c. 
of  alcohol  with  5  c.c.  of  water.  Test  the  conductivity  of  each  solution. 

Classify    the    substances    used    above    according    to    their   electrical 
conductivities. 

2.  Properties    of  Acids.  —  Compare   the   conductivity    of   dilute 
hydrochloric  acid,   dilute   sulphuric   acid,   dilute  acetic    acid,  and  pure 
water.     Compare  the  taste  of  each  of  the  acids  by  stirring  a  few  drops 
into  a  small  beaker  of  water  and  tasting  one  drop  of  the  resulting  solu- 
tion.    Compare  the  action  of  each  on  a  piece  of  zinc,  touching  the  zinc 
with  a  platinum  wire  if  the  acid  alone  has  no  effect.     Compare  the 
effect  of  each  upon  litmus  paper. 

Name  four  components  of  an  acid  solution.  What  is  the  relative 
amount  of  these  components  in  the  different  acid  solutions  examined? 
What  component  is  common  to  all  acids,  and  how  docs  the  strength  of 
an  acid  depend  on  the  amount  of  this'  component  ? 

3.  Properties  of  Bases Dissolve  about   I   gram  of  sodium  hy- 
droxide   in    10    c.c.   of    water.       Dilute    about    2    c.c.  of   concentrated 
ammonium    hydroxide    solution    (sp.  gr.   0.90)   with    10   c.c.  of   water. 


74  Electrolytic  Dissociation   Theory 

Compare  these  two  solutions  with  respect  to  (a)  conductivity,  (b)  effect 
upon  litmus,  (c)  taste  (after  diluting  very  considerably),  (d)  slippery 
feeling  upon  rubbing  a  drop  of  the  solution  between  the  fingers. 

Name  four  components  of  each  solution.  What  is  the  relative 
amount  of  eacJi  component  ?  What  component  is  common  to  the  solu- 
tions of  all  bases,  and  how  does  tJie  strength  of  a  base  depend  upon  the 
amount  of  this  component  f 

4.  Neutralization   of  an   Acid   and   a    Base Run    10   c.c.    of 

a  normal  solution  of  HC1  from  a  burette  into  a  dry  beaker.     Add  one 
drop  of  a  litmus  solution,  and  then  run  in  slowly  from  another  burette 
a  normal  solution  of  NaOH,  until  a  single  drop  just  changes  the  color 
to  blue.      Make  sure  that   one  drop  of  HC1  will  now  cause  the  red 
color  to  reappear. 

Arrange  such  a  form  of  conductivity  vessel  that  the  distance  be- 
tween the  electrodes  can  be  so  adjusted  that,  with  a  normal  solution  of 
hydrochloric  acid,  the  filament  of  the  lamp  will  glow  faintly.  The 
apparatus  described  with  Electrodes  B  serves  well  for  this  purpose. 

With  this  apparatus  compare  the  conductivity  of  both  the  acid  and 
the  base  separately  with  that  of  the  neutralized  solution. 

What  component  of  each  solution  has  disappeared  during  the 
neutralization  ?  What  components  have  not  been  affected  by  the  proc- 
ess ?  Write  the  equation  for  this  reaction.  What  is  the  cause  of  the 
completeness  of  this  reaction  ? 

5.  The  Law  of  Mass  Action.  —  Place  2  c.c.  of  a  saturated  solu- 
tion of  silver  acetate  in  each  of  two  test  tubes.     Add  to  one  tube  a 
small  crystal  of  silver  nitrate  (not  more  than  0.05  gram).     Agitate  the 
solution  until  the  crystal  has  completely  dissolved.     If  no  change  is 
noticed  at  once,  set   the  tube  aside  for  several  minutes  and  observe 
again. 

To  the  second  tube  add  a  small  crystal  of  sodium  acetate  instead  of 
silver  nitrate,  and  observe  as  before. 

Note  I.  Silver  acetate  is  only  slightly  soluble,  that  is,  about  10 
grams  per  liter  at  20°  C.  * 

Note  2.  Put  all  the  waste  from  this  experiment  in  the  bottle 
marked  silver  residues. 

Repeat  the  above  experiment,  substituting  a  saturated  solution  of 


Experiments  75 

potassium  chlorate  for  that  of  silver  acetate.  Dissolve  in  two  portions 
of  this,  respectively,  small  crystals  of  potassium  chloride  and  sodium 
chlorate,  and  observe  as  above. 

Explain  the  application  of  the  Mass  Action  Law  in  these  experi- 
ments. Define  what  is  meant  by  solubility  product. 

6.  Neutralization  of  a  Weak  Acid  and  a  Weak  Base Neu- 
tralize 10  c.c.  of  a  normal  solution  of  acetic  acid  with  a  normal  solution 
of  ammonium  hydroxide.     (See  Experiment  4.) 

Compare  the 'conductivity  of  the  acid  and  the  base  separately  with 
that  of  the  neutralized  solution.  Use  Electrodes  C. 

Compare  the  ionization  of  weak  bases  and  acids  with  that  of  their 
salts,  and  explain  how  during  the  process  of  neutralization  the  con- 
centrations of  the  various  components  cJiange  in  accordance  with  the 
Principle  of  Mass  Action. 

7.  The  Displacement  of  a  Weak  Acid  from  Its  Neutral  Salt  by 

a  Stronger  Acid (a)  To  2  c.c.  of  sodium  benzoate  solution  (50  grams 

per  liter)  add  a  few  drops  of  hydrochloric  acid,  (sp.  gr.  1.12).     Sodium 
benzoate  is  a  strong  electrolyte,  yielding  Na+-  and  (CTH5O2)~-ions. 

Eliminating  all  the  possible  components  which  may  have  been  formed 
with  the  properties  of  which  we  are  already  acquainted,  what  must  be 
the  new  substance  the  formation  of  which  is  made  evident? 

(b)  Heat  2  c.c.  of  sodium  acetate  solution.     Do  you  observe  any 
odor  to  the  solution?     Add  2  c.c.  of  dilute  sulphuric  acid,  (sp.  gr.  1.25) 
and  heat  again.     What  odor  do  you  now  observe  ? 

Explain,  in  accordance  with  the  Principle  of  Mass  Action,  the  forma- 
tion of  the  component  which  is  detected  by  this  odor. 

(c)  To  5  c.c.  of  sodium  carbonate  solution  add  acetic  acid,  a  few 
drops  at  a  time,  till  action  ceases.     What  is  the  gas  formed? 

If  this  gas  were  prevented  from  escaping,  by  keeping  the  solution 
under  a  heavy  pressure,  what  component  would  then  form  which  other- 
wise would  not  form  t^  any  considerable  extent  ?  Write  the  equations 
showing  the  conditions  of  equilibrium  between  the  different  components 
of  this  solution.  How  dots  the  escape  of  the  gas  affect  these  conditions  ? 

(a7)  Treat  a  small  quantity  of  marble  dust  (calcium  carbonate)  with 
some  dilute  hydrochloric  acid,  and  see  if  it  will  all  dissolve. 

Is  calcium  carbonate  appreciably  soluble  in  water  ?     Can  you  explain 


76  Electrolytic  Dissociation  Theory 

the  solvent  action  of  hydrochloric  acid  in  the  same  manner  as  you 
explained  the  action  of  the  acid  in  (c)  ?  If  so,  wJiat  must  be  true  con- 
cerning the  solubility  of  calcium  carbonate  in  pure  water? 

8.  The  Displacement  of  a  Weak  Base  from  Its  Neutral  Salt  by 
Means  of  a  Stronger  Base. —  (a)  Warm  2  c.c.  of  ammonium  chloride 
solution,  and  observe  if  there  is  an  odor.     Add  2  c.c.  of  sodium  hydrox- 
ide solution,  and  again  observe  if  there  is  an  odor. 

What  new  component  is  here  shown  by  the  odor  to  have  been  formed? 
To  how  great  an  extent  was  it  formed?  Explain  ^vhat  further  change 
a  small  fraction  of  this  component  undergoes  to  produce  the  substance 
which  gives  the  odor. 

(b)  To  2  c.c.  of  -magnesium  sulphate  solution  add  a  little  sodium 
hydroxide  solution. 

What  is  the  new  substance  which  is  formed?  Had  this  new  sub- 
stance not  been  insoluble,  to  how  great  an  extent  would  it  have  formed? 

9.  Tests  for  the   Presence   of  Certain   Ions   in  a  Solution.  — 
(a)  Place  some  concentrated  calcium  chloride  solution  in  one  test  tube, 
and  dilute  one  drop  of  it  in  another  with  one-half  of  the  test  tube  full 
of  water.     Add  a  drop  or  two  of  dilute  sulphuric  acid  to  each. 

What  is  the  precipitate  formed  in  one  case  ?  Could  its  formation 
be  used  as  a  test  for  one  of  the  ions  of  calcium  chloride  ?  If  so,  zvould 
it  be  a  delicate  test  ?  What  may  be  said  of  the  solubility  product  of  the 
precipitate  f 

Test  the  electrical  conductivity  of  a  saturated  solution  of  this 
substance. 

(b)  Pour  4  c.c.  of  silver  nitrate  solution  into  a  test  tube,  and  dilute 
with  water  enough  to  half  fill  the  tube.  Then  add  a  solution  of  sodium 
chloride  until  precipitation  is  complete.  Pour  the  contents  of  the  tube 
through  a  filter,  and  wash  the  filter  with  pure  water  until  I  c.c.  of  the 
wash  water  gives  no  turbidity  with  a  drop  of  silver  nitrate.  Now 
break  the  point  of  the  filter  with  a  glass  rod,  and  wash  the  precipitate 
into  a  3-inch  test  tube  together  with  4  c.c.  of  water.  With  the  thumb 
over  the  mouth  of  the  tube,  shake  thoroughly  for  some  moments ;  then, 
without  filtering,  test  the  electrical  conductivity  of  the  liquid. 

To  one  test  tube  half  full  of  water  add  one  drop  of  silver  nitrate ; 
to  another  add  one  drop  of  sodium  chloride  solution.  Pour  the  contents 
of  the  two  tubes  together. 


Experiments  77 

What  is  t/ie  new  substance  formed  in  this  experiment  ?  F*or  what 
ion  is  its  formation  a  test  ?  Is  this  a  delicate  test  ?  How  does  the 
solubility  product  of  this  substance  compare  with  that  of  calcium 
sulphate  ?  WJiat  components  were  contained  in  the  liquid  which  ran 
through  the  filter,  and  in  what  relative  amounts  ?  If  this  liquid  had 
been  evaporated^  what  would  have  been  the  nature  of  the  residue? 
Which  components  disappeared  during  the  evaporation,  and  which 
increased  in  amount  ? 

(c)  Add  a  drop  of  silver  nitrate  solution  to  3  c.c.  of  each  of  the 
following  solutions :  calcium  chloride,  cobalt  chloride,  copper  chloride. 
Add  a  drop  of  silver  sulphate  solution  to  some  sodium  chloride  solution. 

What  ion  is  common  to  all  aqueous  solutions  of  metallic  chlorides  ? 
Of  simple  silver  salts  ? 

(d)  Mix   i   c.c.  of  alcohol  with  3  c.c,  of  water.     Add  three  drops  of 
chloroform,  and  mix  by  shaking.     Test  the  conductivity  of  this  solution, 
and  then  add  a  drop  of  silver  nitrate  solution. 

Account  for  the  different  behavior  of  chlorine  in  this  compound  and 
in  such  compounds  as  sodium  chloride. 

10.  Simple  and  Complex  Ions (a)  Add  a  few  drops  of  lead 

nitrate  solution  to  3  c.c.  of  solutions  of  (i)  sodium  sulphide,  (2)  sodium 
sulphate,  (3)  sodium  thiosulphate. 

Account  for  the  different  behavior  of  sulphur  in  these  three  com- 
pounds. 

(b)  Add  a  drop  of  silver  nitrate  to  solutions  of  potassium  chlorate 
and  chloracetic  acid,  (HC2H2C1O2). 

(c)  Add  5  c.c,  of  ammonium  hydroxide  (sp.  gr.  0.96)  to  i  c.c.  of 
dilute  silver  nitrate  solution,  and  then  add  some  potassium  chloride 
solution. 

What  has  become  of  the  silver  ions  ? 

11.  The  Power  of  Solvents  Other  than  Water  to  Cause  Elec- 
trolytic Dissociation (a)  Test  a  solution  of  hydrogen  chloride  in 

toluene  with  regard  to  (i)  its  conductivity;  (2)  its  action  upon  a  bit  of 
marble ;  (3)  its  action  upon  a  piece  of  zinc.     Touch  the  zinc  beneath 
the  liquid  with  a  platinum  wire.     All  apparatus  and  materials  used  in 
this  experiment  must  be  perfectly  free  from  moisture. 

What  are  the  components  of  tJiis  solution  ?  How  does  a  solution  of 
hydrogen  chloride  in  toluene  differ  from  one  in  water? 


78  Electrolytic  Dissociation   Theory 

(b)  To  2  c.c.  of  the  above  solution  in  a  small  test  tube  add  2  c.c.  of 
water,  and  shake.     Now  insert  the  electrodes  till  they  dip  in  the  aqueous 
layer. 

(c)  Test  a  solution  of  hydrogen  chloride  in  absolute  alcohol  as  in  (a). 

12.  The    Relative    Electrolytic    Solution    Tension    of    Differ- 
ent Metals.  —  (a)  Place  some  scraps  of  zinc  in  5  c.c.  of  dilute  copper 
sulphate  solution.     Allow  them  to  stand,  with  frequent  shaking  (why  ?), 
until  the  chemical  change  is  complete.     Pour  off  a  little  of  the  clear 
liquid  and  add  a  little  ammonia  water  to  it.     Add  also  ammonia  water 
to  some  copper  sulphate  solution.     Add  some  ammonium  sulphide  to 
another  small  portion  of  the  liquid,  and  also  to   some  zinc   sulphate 
solution. 

What  ions  are  shown  to  be  present  or  absent  by  the  tests  with  am- 
monia water  and  with  ammonium  sulphide  ?  What  has  caused  the 
above  change? 

(&)  Put  some  pieces  of  copper  wire  in  2  c.c.  of  silver  nitrate  solu- 
tion. After  standing  a  few  minutes  test  the  liquid  for  silver  ions. 

What  can  be  said  of  the  relative  electrolytic  solution  tension  of  the 
metals  studied  in  (a)  and  (b]  ?  Judging  from  the  behavior  of  these  metals 
with  acids,  where  would  you  place  hydrogen  in  this  classification  f 

13.  The  Relative  Tendency  of  Non-metallic  Elements  to  Pass 

into  the  Ionic  Form (a)  Add  bromine  water  and  iodine   solution 

to  separate  portions  of  hydrogen  sulphide  water. 

(b)  What  is  the  effect  of  a  solution  of  each  of  the  free  halogens 
upon  solutions  of  the  potassium  or  hydrogen  compounds  of  the  other 
halogens  ?      (If  you  do  not  know  this,  try  the  necessary  experiments 
with  solutions  to  be  found  in  the  laboratory.)     To  test  for  the  presence 
of  free  bromine  or  iodine  in  a  solution  add  about   i  c.c.  of  carbon  bi- 
sulphide to  the  liquid  ;  shake,  and  allow  the  globule  of  carbon  bisulphide 
to  settle.     This  will  have  collected  the  free  halogen,  imparting  to  it, 
if  bromine,  a  red  color ;  if  iodine,  a  violet  color. 

(c)  What  substance  separates  from  a  solution  of  hydrogen  sulphide 
which  has  been  allowed  to  stand  where  it  could  absorb  oxygen,  as  in 
a  bottle  the  stopper  of  which  is  frequently  removed  ? 

Arrange  the  halogens,  sulphur,  and  oxygen,  as  nearly  as  you  can  tell 
from  the  data  in  (a},  (b),  (c),  in  the  order  of  their  tendency  to  take  the 
form  of  negative  ions. 


Experiments  79 

14.  Hydrolysis.  —  (a)  Dissolve  about  half  a  gram  each  of  zinc 
sulphate,  sodium  chloride,  and  sodium  carbonate  in  a  little  water,  and 
test  each  solution  with  red  and  blue  litmus. 

Explain  the  relation  of  hydrolysis  to  the  observed  results. 

(b)  Place  i  c.c.  of  a  concentrated  solution  of  zinc  chloride  in  a  test 
tube.  Fill  the  tube  half  full  with  water.  Now  add  hydrochloric  acid 
drop  by  drop  till  the  precipitate  disappears. 

OTT 

Assuming  tJie  precipitate  to  have  been  basic  zinc  chloride,   Zn<(        , 

explain  its  formation  from  the  components  of  the  solution,  and  explain 
why  a  slight  excess  of  hydrochloric  acid  causes  its  disappearance. 


15.  Saturate  5   c.c.  of  a   lo-per  cent,  zinc  chloride   solution  with 
hydrogen  sulphide.      Filter  off  the  precipitate,  and  add  to  the  filtrate 
some  ammonium  sulphide. 

To  a  second  portion  of  zinc  chloride  solution  add  i  c.c.  of  dilute 
hydrochloric  acid  (sp.  gr.  1.12),  and  again  saturate  with  hydrogen 
sulphide. 

To  a  third  portion  of  zinc  chloride  solution  add  2  c.c.  of  concentrated 
sodium  acetate  solution,  saturate  with  hydrogen  sulphide,  filter,  and  to 
the  filtrate  add  ammonium  sulphide. 

Explain  how  the  degree  to  which  hydrogen  sulphide  can  dissociate 
is  affected  by  the  components  present  in  tJie  three  cases  above,  and  wJiat 
effect  the  degree  of  this  dissociation  Jias  upon  the  formation  of  a 
precipitate. 

1 6.  Prepare  four  samples  of  sulphuric  acid  of  different  strengths,  as 
follows :  — 

1.  Dilute   i  volume  of  acid  of   1.25   sp.  gr.  with  an  equal  volume 
of  water. 

2.  To  2  volumes  of  acid  of   1.25  sp.  gr.  add  i  volume  of  concen- 
trated acid,  sp.  gr.  1.84. 

3.  To  I  volume  of  water  add  4  volumes  of  acid,  sp.  gr.  1.84. 

4.  To  3  volumes  of  acid,  1.84  sp.  gr.,  add  i  volume  of  fuming  sul- 
phuric acid.     This  yields  a  liquid  which  contains  a  slight  excess  of  SO3. 

Test  the  conductivity  of  each  of  these  liquids,  using  Electrodes  B. 


8o  Electrolytic  Dissociation   Theory 

Treat  each  with  feathered  zinc,  warming  after  observing  whether  any 
action  takes  place  in  the  cold.  Test  the  gases  evolved  in  each  case  for 
hydrogen,  sulphur  dioxide,  and  hydrogen  sulphide,  and  observe  also  if 
any  sulphur  collects  on  the  sides  of  the  tube. 

Repeat  the  above  experiments,  using  copper  turnings  instead  of 
feathered  zinc. 

Draw  conclusions  from  the  conductivity  and  from  the  products 
formed  in  each  case  as  to  what  must  have  been  the  components  of  the 
liquid,  and  explain  how  the  relative  electrolytic  solution  tensions  of 
zinc,  hydrogen,  and  copper  affect  the  result. 


APPENDIX 

DEGREE    OF    DISSOCIATION    OF    SOME    OF    THE    MOST    IMPORTANT 

ELECTROLYTES 

SALTS 

NEUTRA^  salts,  with  very  few  exceptions,  are  highly  dissociated. 
If  they  are  classified,  as  in  the  table  below,  according  to  the  valence 
of  their  ions,  it  is  found  that  all  belonging  to  any  one  class  have  prac- 
tically the  same  degree  of  dissociation.  The  four  classes  which  are 
indicated  may  be  typified  by  KNO3,  Ba(NO3)2,  K2SO4,  and  ZnSO4, 
respectively. 


Type  of  salt. 

Percentage  dissociation  in  0.1 
equivalent  solution. 

R+R-     

86 

R++R2-    

72 

R2+R--     

72 

R++R--       

45 

ACIDS 


Substance. 

Percentage  dissociation  in  0.1 
equivalent  solution. 

HC1,  HBr,  HI,  HNO3      | 

90 

HC1O3,  HC1O4,  HMnO4  \    
H2SO4  

60 

H2C2O4     

QJ. 

H2SO3  

20 

H8PO4  

13 

H8AsO4     

H 

HF     

9 

HC2H3O*     

1  4. 

H2CO3       

010 

H2S    

0  0^ 

HCN     

0  01 

s. 


82 


Appendix 
BASES 


Substance, 


KOH,  NaOH 
Ba(OH)2    .    . 
NH4OH    .    . 


Percentage  dissociation  in  0.1 
equivalent  solution. 


86 
75 
1.4 


Ca(OH)2,  Mg(OH)2  are  but  slightly  soluble,  but  so  far  as  they  do 
dissolve  they  are  dissociated  to  about  the  same  extent  as  Ba(OH)2  in 
a  solution  of  the  same  concentration. 

The  hydroxides  of  the  heavy  metals  are  very  insoluble  and,  as  a 
rule,  very  weakly  basic. 

AgOH  is  soluble  to  the  extent  of  one  part  in  15,000  of  water,  in 
which  solution  about  33  per  cent,  of  its  molecules  are  ionized.  It  is 
thus  a  moderately  strong  base. 

Hydroxides  of  the  type  Zn(OH)2,  Fe(OH)2,  Mn(OH)2  are  much  less 
basic  than  AgOH,  while  hydroxides  of  the  type  Fe(OH)3,  Cr(OH)3, 
A1(OH)3  are  least  basic  of  all. 

The  dissociation  of  pure  water  into  H+-  and  OH  "-ions  amounts  to 

of   l  Per  cent' 


INDEX 


Acids,  properties  of 33>  73 

ionization  of 33»  &1 

Ammonium  chloride,  vaporization  of    .    .    9 

Apparatus  for  experiments 71 

Arrhenius I 

Avogadro's  Theory 7>  9 

Bases,  properties  of 34»  73 

ionization  of 34 >  &2 

Boiling  point,  raising  of  the 5 

Boyle's  Law 6 

Carbonates,  solubility  in  acids 39 

Charles'  Law 7 

Chemical  activity 10 

greater    with    electro- 
lytes   10 

Chlorine,  reduction  of 52 

Chloroform,  behavior  as  non-electrolyte  .  10 

Chromate  ions,  reduction  of 55 

Conductors,  definition  of I 

electrolytic I 

metallic I 

Copper  sulphide,  precipitation  of  .    .    .    .40 
Current,  definition  of 12 

Effect  of  mass,  the 21 

Electrical  charges  upon  the  ions    .    .    .    .  16 

Electrodes,  description  of 72 

Electrolysis,  definition  of i 

description  of 13 

of  copper  sulphate  ....  17 
of  hydrochloric  acid  .  n,  13, 17 
of  potassium  sulphate  .  18,  50 

Electrolytes,  definition  of .     i 

Electrolytic  conduction u,  73 

Electrolytic    Dissociation   Theory,   state- 
ment of i 

Electrolytic  solution  pressure    .   .  45,  46,  78 


Equilibrium,  definition  of 22 

Experiments 71 

Ferrous  sulphide,  precipitation  of    ...  40 
Freezing  point,  lowering  of  the 2 

Gas  pressure,  discussion  of 6 

of  mol  of  any  gas    ....    7 

Hydrolysis 42,  79 

of  aluminum  sulphide    ....  44 

of  ferric  chloride 44 

of  potassium  cyanide    ....  43 
of  salts    from  weak  acid  and 
base 44 

Insoluble  compounds,  formation  of  .    .    .26 

Ionization,  cause  of  ....        12 

changes  with  concentration    .  26 

degree  of 25 

Ionization  pressure 49,  78 

Ions,  characteristic  reactions  of  .    .  31,  57,  76 

complex 32,  77 

definition  of i,  12 

electrical  charges  upon 16 

magnitude  of  charges  upon    .    .    .    .  18 

movement  of 14 

simple .   31,  77 

Ions  of  sulphur,  oxidation  and  reduction 

of 54 

Iron,  oxidation  of 51 

Law  of  Boyle 6 

of  Charles 7 

of  Mass  Action 22,  74 

of  Mass  Action,  application  of,  with 

complex  ions 33 

of  Mass  Action,  demonstration  of  .  23 

Lowering  of  the  freezing  point,  molecular  .    3 

the  .    2 


Index 


Magnitude  of  charges  upon  the  ions     .    .18 

Mass  Action,  law  of 22,  74 

Membranes,  osmotic 8 

semi-permeable 8 

Mercury,  ions,  reduction  of 53 

Mol,  definition  of 3 

Molecular  lowering  of  the  freezing  point  .    3 
Molecular  raising  of  the  boiling  point  of 

water 5 

Movement  of  the  ions 14 

demonstration   of   .  15 
rate  of 15 

Neutralization 35»  74 

heat  of 35 

of  weak  acid  or  base  .  36,  75 

Non-conductors i 

Non-electrolyte,  definition  of 3 

Number  of  molecules  in  solution,  deter- 
mined by  boiling  point     ....    5,  19,  20 
by  freezing  point  .    .    .    .    4,  19,  20 
by  osmotic  pressure    .    .    9,  19,  20 

Osmotic  pressure 6,  46 

definition  of 7 

measurement  of  ....  8 

Oxidation,  definition  of 5r>52>53 

of  iron 51 

Permanganate  ions,  reduction  of   ....  56 
Polarity  in  chemical  compounds     .    .    .    .  1 1 
Properties  of  solutions  associated  with 
electrical  conductivity 2 

Raising  of  the  boiling  point -5 

Reactions  at  the  electrodes 17 


Reactions  of  ions  —  metals     ....    58-67 
non-metals      .    .    67-70 

Reduction,  definition  of 51,  53 

of  chlorine 52 

Reversible  reactions 21,  25 

Salt  of  weak  acid,  action  of  strong  acid 

uP°n 37,75 

Salt  of  weak  base,  action  of  strong  base 

upon 41,  76 

Salts,  formation  of 34 

ionization  of 35,  37,  81 

Saturated  solutions 28 

addition   of    other 

salts  to 29 

Solubility  product 28,  29,  74 

applications  in  chem- 
ical analysis   .    .    .    .32 

Solution  pressure  .    ......      45,46,78 

apparent  exceptions   .    .  48 
of  hydrogen      .    .    .    47,  48 

of  metals      47 

Sulphides,  metallic,  precipitation  of  .    .    .40 
Summary  of  Chapter  I 19 

Tin,  ions,  oxidation  of 53 

Volatile  products,  effect  of  formation  of,    38 

Water,  dissociating  power  of  .  .  .  .  13,77 
influence  of  ionization  of  .  .  42,  50 
Weak  acids,  effect  of  salts  with  common 

ion  up6n 41 

Weak  bases,  effect  of  salts  with  common 

ion  upon 41 

Zinc  sulphide,  precipitation  of 40 


QD-T6 


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